UC-NRLF 


GIFT  OF 
Hheology  Seminar 


EM31NEER1NG  L.~ 


CONDUCTIVITIES  AND  VISCOSITIES  IN  PURE  AND 
IN  MIXED  SOLVENTS 


RADIOMETRIC  MEASUREMENTS  OF  THE 

IONIZATION  CONSTANTS  OF 

INDICATORS 


BY 


HARRY  C.  JONES  AND  COLLABORATORS 


WASHINGTON,  D.  C. 

PUBLISHED  DY  THE  CARNEGIE  INSTITUTION  OF  WASHINGTON 
1915 


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CONDUCTIVITIES  AND  VISCOSITIES  IN  PUEE  AND 
IN  MIXED  SOLVENTS 


RADIOMETRIC  MEASUREMENTS  OF  THE 

IONIZATION  CONSTANTS  OF 

INDICATORS,  ETC. 


BY 


HARRY  C.  JONES  AND  COLLABORATORS 


WASHINGTON,  D.  C. 

PUBLISHED  BY  THE  CARNEGIE  INSTITUTION  OF  WASHINGTON 
1915 


"  • 
Library 


CARNEGIE  INSTITUTION  OF  WASHINGTON 

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PKESS    OF   GIBSON    BKOTHEKS 
WASHINGTON,  D.  C. 


PREFACE. 

The  investigations  discussed  in  this  monograph  might  seem,  at  first 
sight,  to  deal  with  a  variety  of  subjects,  some  of  which  have  no  very  close 
relation  to  the  others;  but  nearly  all  of  these  lines  of  work  were  sug- 
gested by  the  solvate  theory  of  solution  as  proposed  here  about  fifteen 
years  ago,  and  have  been  carried  out  with  the  specific  purpose  of 
ascertaining  their  bearing  on  this  theory. 

Dr.  Davis  studied  the  viscosities  of  solutions  of  caesium  salts  in  mixed 
solvents,  and  in  mixtures  of  the  associated  liquids  water,  formic  acid, 
and  acetic  acid,  in  its  relation  to  the  theory  of  viscosity  proposed  by 
.Tones  several  years  ago. 

Dr.  Davis  and  Dr.  Putnam  investigated  the  dissociating  power  of 
f ormamid  as  a  solvent,  in  connection  with  its  bearing  on  a  theory  of  the 
relation  between  the  dissociating  power  of  a  solvent  and  other  physical 
properties  of  that  liquid;  namely,  its  dielectric  constant  and  its  own 
association. 

Dr.  Shaeffer  and  Dr.  Paulus  used  the  new  spectroscope  constructed 
to  study  the  nature  of  "solvent  bands,"  to  determine  the  constants  of 
indicators — a  quantity  absolutely  essential  to  the  scientific  use  of  these 
substances  in  quantitative  analysis. 

This  work  was  extended  by  Dr.  Paulus  and  Mr.  Hutchinson  to  the 
indicator,  corallin. 

Having  found  such  a  marked  difference  between  the  physical  behavior 
towards  light  of  free  and  of  combined  water,  Dr.  J.  E.  L.  Holmes  took 
up  the  study  of  a  reaction  effected  by  the  ions  of  water — the  saponifica- 
tion  of  an  ester — to  see  whether  any  difference  in  the  chemical  activity 
of  free  and  of  combined  water  could  be  detected.  It  seems  that  such 
a  difference  exists,  combined  water  having  apparently  greater  chemical 
activity  than  free  water. 

This  work  was  extended  by  Mr.  Connolly  to  another  reaction  which 
proceeds  much  more  rapidly — the  hydration  of  acetic  anhydride. 
This  work  has  not  yet  been  extended  sufficiently  to  justify  any  final 
conclusion. 

The  work  begun  two  years  ago  on  the  conductivities  of  organic  acids 
in  ethyl  alcohol,  was  greatly  extended  by  Dr.  Lloyd  and  Mr.  Wiesel, 
comprising  the  study  of  nearly  forty  of  the  more  common  organic  acids 
in  this  solvent.  The  aim  of  this  work  is  to  arrive  at  a  knowledge  of  the 
dissociation  of  these  substances  by  alcohol,  and  the  degree  of  solvation 
of  these  compounds  in  alcohol. 

Dr.  Watkins  has  extended  work  begun  in  this  laboratory  about  a 
dozen  years  ago  on  the  conductivity  and  dissociation  of  salts  by  water. 
He  has  studied  a  number  of  the  less  common  salts,  and  has  obtained 
very  accurate  results  which,  in  general,  confirm  the  conclusions  drawn 
from  our  earlier  work  in  this  field. 


iv  Preface. 

The  different  physical  behavior  of  free  and  of  combined  water 
suggested  to  Dr.  Shaeffer  the  desirability  of  studying  the  relative  disso- 
ciating powers  of  free  and  of  combined  water.  This  work  was  under- 
taken by  Mr.  Ordeman.  He  prepared  isochloric  solutions  of  the 
slightly  hydrated  salt  potassium  chloride,  and  of  the  strongly  hydrated 
salt  calcium  chloride,  and  used  these  solutions  as  solvents  in  which  to 
dissolve  chlorides  and  measure  their  dissociations.  The  dissociation 
of  a  number  of  salts  in  these  solvents  was  measured;  and  while  we  are 
not  yet  prepared  to  draw  any  final  conclusions,  it  seems  probable  that 
"combined"  water  has  rather  less  dissociating  power  than  "free" 
water. 

The  absorption  and  adsorption  of  the  slightly  hydrated  salt  potas- 
sium chloride  by  soils,  is  a  problem  of  both  scientific  and  technical 
importance.  Work  has  been  in  progress  in  my  laboratory  on  this 
problem  during  the  past  year  by  Mr.  McCall,  with  the  cooperation  of 
Messrs.  Hildebrandt,  Johnston,  F.  S.  Holmes,  and  Trelease.  They  find 
that  a  part  of  the  potassium  chloride  is  absorbed  or  combined  chem- 
ically with  the  soil  particles,  and  a  part  is  adsorbed  or  in  a  state  of 
physical  union  with  the  particles  of  the  soil.  Under  certain  conditions 
the  solvent  may  be  more  rapidly  adsorbed  than  the  salt. 

The  results  of  these  investigations,  many  of  which  have  been  carried 
out  with  the  aid  of  grants  generously  awarded  me  by  the  Carnegie 
Institution  of  Washington,  are  all  recorded  in  this  publication. 

HARRY  C.  JONES. 

JOHNS  HOPKINS  UNIVERSITY, 

June,  1915. 


CONTENTS. 


CHAPTER  I. 

PAGE 

THE  VISCOSITIES  OF  SOLUTIONS  OF  C.-ESIUM 
SALTS  IN  MIXED  SOLVENTS;  THE  VISCOS- 
ITIES OF  BINARY  MIXTURES  OF  THE 
ASSOCIATED  LIQUIDS,  WATER,  FORMIC 
Aero  AND  ACETIC  ACID,  TOGETHER  WITH 
SOME  NEW  FORMS  OF  APPARATUS. 

Introduction 3 

Historical 3 

Apparatus  and  Procedure (> 

Discussion  of  Results 8 

The  Viscosities  of  Binary  Mixtures  of  the 
Associated  Liquids,  Water,  Formic 

Acid  and  Acetic  Acid 10 

Some  new  Forms  of  Apparatus: 

I,  a  Substitute  for  the  Twin-Bulb  Trap  in 
Toluol-Mercury  Thermo-regulators . ...  13 

II,  a  new  Form  of  Pyknometer  for  Liquids  14 

CHAPTER  II. 

THE  CONDUCTIVITY  AND  VISCOSITY  OF  SOLU- 
TIONS OF  CERTAIN  SALTS  IN  FORMAMID 
AS  A  SOLVENT. 

Previous  Work  in  Formamid 17 

Formamid  as  a  Solvent 18 

Purification  of  the  Solvent 20 

Salts 23 

Solutions 23 

Conductivity  Apparatus 24 

Viscosity  Apparatus 24 

Thermostats 24 

Discussion  of  Results 24 

Comparison  of  the  Various  Solvents 24 

Sodium  Bromide  in  Formamid 25 

Sodium  Iodide  in  Formamid 25 

Sodium  Chromate  in  Formamid 26 

Potassium  Chloride  in  Formamid 26 

Potassium  Iodide  in  Formamid 27 

Potassium  Sulphocyanate  in  Formamid ....  27 

Ammonium  Bromide  in  Formamid 28 

Ammonium  Iodide  in  Formamid 28 

Tetramethylammonium  Chloride  in  Forma- 
mid    29 

Tetraethylammonium  Iodide  in  Formamid. .  29 

Rubidium  Chloride  in  Formamid 30 

Rubidium  Bromide  in  Formamid 30 

Rubidium  Iodide  in  Formamid 31 

Rubidium  Nitrate  in  Formamid 31 

Casium  Chloride  in  Formamid ...  32 


CHAPTER  II — Continued. 

PAGE 

Cassium  Nitrate  in  Formamid 32 

Lithium  Nitrate  in  Formamid 33 

Barium  Chloride  in  Formamid 33 

Mercuric  Chloride  in  Formamid 34 

Cobalt  Bromide  in  Formamid 34 

Sodium  Chromate 37 

Mercuric  Chloride 38 

Cobalt  Bromide 38 

Csesium  Nitrate  and  Chloride 38 

Summary  or  Results 39 

CHAPTER  III. 

RADIOMETHIC  MEASUREMENTS  OF  THE  IONI- 
ZATION  CONSTANTS  OF  INDICATORS  I. 

Purpose  of  the  Investigation 41 

Historical 42 

The  R  adiomicrometer 44 

The  Spectroscope 48 

Source  of  Light 50 

The  Cells 50 

Arrangement  of  Apparatus 51 

The  Differential  Method 52 

Theoretical  Discussion 52 

Experimental  Work  on  Methyl  Orange 57 

Results  with  Methyl  Orange 59 

Phenolphthalein 64 

Results  with  Phenolphthalein 66 

Summary 72 

CHAPTER  IV. 

RADIOMETRIC  MEASUREMENTS  OF  THE  IONI- 
ZATION  CONSTANTS  OF  INDICATORS  II. 

Theoretical  Discussion 73 

Preliminary  Work  on  Rosolic  Acid 75 

Results  with  Rosolic  Acid 79 

Rosolic  Acid  as  a  Dibasic  Acid 80 

Summary 82 

CHAPTER  V. 

THE  ACTION  OF  SALTS  WITH  WATER  OF 
HYDRATION  AND  WITHOUT  WATER  OF 
HYDRATION  ON  THE  VELOCITIES  OF 
SAPONIFICATION  OF  ESTERS. 

Reaction  Chosen 85 

Historical 86 

Hydrolysis 88 


VI 


Contents. 


CHAPTER  V — Continued. 

PAGE 

Statement  of  Problem 90 

Apparatus 90 

Solutions 90 

The  Esters 91 

The  Base  and  Indicator 91 

Method  of  Procedure 93 

Calculations 94 

Data 94 

Discussion  of  Results 106 

Summary 110 

CHAPTER  VI. 

THE  EFFECT  OF  NEUTRAL  SALTS  ON  THE 
HYDRATION  OF  ACETIC  ANHYDRIDE. 

Hydration  of  Acetic  Anhydride 112 

Neutral  Salt  Action 113 

Statement  of  the  Problem 114 

Purification  of  Materials 115 

Apparatus  and  Solutions 115 

Manipulation 116 

Discussion 118 

Conclusion .  118 


CHAPTER  VII. 

CONDUCTIVITY  OF  CERTAIN  ORGANIC  ACIDS 
IN  ETHYL  ALCOHOL. 

Historical  Sketch 119 

Experimental  Work 129 

Reagents 129 

Apparatus 130 

Procedure 131 

Results 133 

Phenylacetic  Acid 134 

Oxyisobutyric  Acid 134 

Brompalmitic  Acid 134 

Malonic  Acid 134 

Ethylmalonie  Acid 134 

Diethyhnalonic  Acid 134 

Propylmalonic  Acid 134 

Dipropylmalonic  Acid 135 

Butylmalonic  Acid 135 

Allylmalonic  Acid 135 

Benzylmalonic  Acid 135 

Monobromsuccinic  Acid 135 

Dibromsuccinic  Acid 135 

Sebasic  Acid 135 

Thiodiglycolic  Acid 135 

Benzilic  Acid 135 

Maleic  Acid 136 

Fumaric  Acid 136 

Itaconic  Acid 136 

Mesaconic  Acid 136 

Phenylpropiolic  Acid 136 

Aconitic  Acid .  136 


CHAPTER  VII — Continued. 

['AGE 

Benzoic  Acid 136 

w-Chlorbenzoic  Acid 137 

w-Nitrobenzoic  Acid 1 37 

Dinitrobenzoic  Add 137 

|   Picric  Acid 137 

Sulphosalicylic  Acid 137 

o-Aminobenzoic  Acid 137 

/>-Aminobenzoic  Acid 137 

o-Toluic  Acid 137 

?>Toluic  Acid I3S 

Cinnamic  Acid 1 38 

Phthalic  Acid 138 

Dichlorphthalic  Acid 138 

Anisic  Acid 138 

Mandelic  Acid 138 

Camphoric  Acid 1 38 

Discussion  of  the  Results 1311 

Relation  between  Composition  and  Conduc- 
tivity   142 


CHAPTER  VIII. 

CONDUCTIVITY,  TEMPERATURE  COEFFI- 
CIENTS OF  CONDUCTIVITY,  AND  PER- 
CENTAGE DISSOCIATIONS  OF  SOME 
RATHER  UNUSUAL  SALTS  IN  AQUEOUS 
SOLUTIONS. 

Apparatus 1 47 

Bridge  and  Rheostat 1 47 

Cells 147 

Constant  Temperature  of  Baths 147 

Containing  Vessels 147 

Solutions 147 

Water 147 

Salts 148 

Standardization 148 

Cell  Constants 148 

Precautions 148 

Results 149 

Sodium  Bromate 150 

Sodium  Sulphocyanate 150 

Sodium  Thiosulphate 150 

Sodium  Dithionate 151 

Sodium  Pyrophosphate 151 

Trisodium  Phosphate 151 

Sodium  Dihydrogen  Phosphate 153 

Sodium  Tungstate 153 

Sodium  Formate 153 

Sodium  Chromate 154 

Sodium  Dichromate 154 

Potassium  Ferricyanide 154 

Ammonium  lodate 155 

Ammonium  Dihydrogen  Phosphate 156 

Ammonium  Chromate ;v.  .  156 

Ammonium  Sulphocyanate 157 

Lithium  Chromate 157 

Rubidium  Iodide .  157 


Contents. 


VII 


CHAPTER  VIII — Continued. 

PAGE 

Discussion  of  Results 157 

Conductivities 157 

Dissociations 158 

Temperature  Coefficients 159 

Summary 160 

CHAPTER  IX. 

THE  DISSOCIATING  POWERS  OF  FREE  AND  OF 
COMBINED  WATER. 

Introduction 161 

Experimental  Apparatus 102 

Water 162 

Isoohloric  Solutions 102 


CHAPTER  IX — Continued. 

PAOB 

Salts 164 

Solutions 164 

Procedure 164 

Discussion  of  Results 166 

CHAPTER  X. 

THE  ABSORPTION  OF  POTASSIUM  FROM 
AQUEOUS  SOLUTIONS  OF  POTASSIUM 
CHLORIDE. 

Resum6  of  Previous  Work 167 

Description  of  Apparatus 170 

Results.  .  .171 


CONDUCTIVITIES  AND  VISCOSITIES  IN  PURE  AND 
IN  MIXED  SOLVENTS 


RADIOMETRIC  MEASUREMENTS  OF  THE  IONIZATION 
CONSTANTS  OF  INDICATORS 


BY 


HARRY  C.  JONES  AND  COLLABORATORS 


CHAPTER  I. 

THE  VISCOSITIES  OF  SOLUTIONS  OF  CESIUM  SALTS  IN  MIXED  SOL- 
VENTS; THE  VISCOSITIES  OF  BINARY  MIXTURES  OF  THE  ASSO- 
CIATED LIQUIDS.  WATER,  FORMIC  ACID.  AND  ACETIC  ACID; 
TOGETHER  WITH  SOME  NEW  FORMS  OF  APPARATUS. 


BY  P.  B.  DAVIS. 

INTRODUCTION. 

We  have  endeavored  for  some  time  to  secure  enough  caesium  salts 
to  study  their  viscosities  in  pure  and  in  mixed  solvents,  but  only  within 
the  past  year  have  we  been  successful.  Through  the  courtesy  and 
cooperation  of  Professor  James  Lewis  Howe  of  Washington  and  Lee 
University,  a  quantity  of  caesium  sulphate  was  placed  at  our  disposal. 
This  was  converted  first  into  the  hydroxide,  then  into  the  carbonate, 
and  finally  into  the  chloride  and  nitrate;  and  with  these  salts  this 
investigation  has  been  carried  out. 

Caesium  is  the  most  electro-positive  of  all  the  elements,  and  is  further 
distinguished  by  possessing  the  largest  atomic  volume,  being  followed 
in  this  respect  by  rubidium  and  potassium,  respectively.  Since  salts 
of  the  latter  two  elements  are  of  great  interest  from  the  viscosity  stand- 
point, it  would  be  expected  that  caesium  salts  would  possess,  to  a  much 
more  pronounced  degree,  any  peculiarities  shown  by  salts  of  rubidium 
and  potassium. 

An  examination  of  the  literature  bearing  on  viscosity  shows  that,  in 
general,  only  the  salts  of  the  three  metals  mentioned  above  are  known 
to  lower  the  viscosity  of  water.  The  effect  of  potassium  and  rubidium 
salts  on  the  viscosities  of  solvents  other  than  water,  and  of  mixtures  of 
such  solvents  with  one  another  and  with  water,  has  been  the  subject  of 
earlier  investigations  in  this  laboratory;  and  this  series  of  investiga- 
tions can  now  be  regarded  as  partially  completed  by  this  study  of 
caesium  salts  in  these  solvents.  The  present  investigation,  therefore, 
has  been  made  to  comprise  a  study  of  the  viscosities  of  the  two  caesium 
salts,  chloride  and  nitrate,  in  water  and  in  binary  mixtures  of  methyl 
alcohol,  ethyl  alcohol,  and  acetone  with  water. 

The  results  obtained  with  these  salts  of  caesium  in  formamid  as  a 
solvent  are  published  in  Chapter  II,  and  further  determinations  of  their 
behavior  in  glycerol  and  glycerol-water  mixtures,  as  well  as  in  mixed 
solvents  containing  formamid  instead  of  water  are  now  in  progress. 

HISTORICAL. 

Jones  and  Lindsay,  in  their  work  on  binary  mixtures  of  the  alcohols 
with  water,  found  that  the  molecular  conductivities  of  solutions  of  salts 
in  these  solvents  were  in  every  case  less  than  the  averages  calculated 
from  the  conductivities  in  the  component  solvents  themselves.  These 

3 


4  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

results,  interpreted  in  the  light  of  the  work  of  Dutoit  and  Aston,  point 
to  a  lower  degree  of  association  in  the  case  of  the  mixed  solvents  than 
in  that  of  the  pure  solvents  themselves.  That  this  is  true  was  shown 
by  Jones  and  Murray,  who  carried  out  cryoscopic  measurements  with 
mixtures  of  water,  formic  acid,  and  acetic  acid;  and  showed  that  the 
molecular  weights  of  these  substances  were  always  less  in  solution  than 
the  molecular  weights  of  the  pure  solvents  themselves,  as  determined 
by  the  method  of  Ramsay  and  Shields. 

Jones  and  Carroll,  in  extending  this  investigation,  showed  that  the 
change  in  the  viscosity  of  the  medium  was  an  important  factor  condi- 
tioning the  decrease  in  the  conductivity  of  salts  in  such  binary  mixtures 
of  liquids,  the  decrease  in  ionization  due  to  changes  in  association  being 
only  one  of  the  factors  to  be  taken  into  account. 

However,  it  remained  for  Jones  and  Veazey  to  offer  a  satisfactory 
explanation  of  the  increase  in  viscosity  which  occurs  when  two  highly 
associated  liquids  are  mixed,  and  also  to  account  for  the  phenomenon 
of  negative  viscosity,  or  the  lowering  of  the  viscosity  of  the  solvent  by 
a  dissolved  substance.  They  carried  out  an  extensive  investigation  in 
mixtures  of  the  alcohols  and  of  acetone  with  water,  and  found  that  the 
conductivity  minima  referred  to  above  were  more  general  than  had 
hitherto  been  supposed,  and  that  this  was  very  closely  related  to  the 
fluidity  minima  observed  in  these  same  mixtures. 

The  work  of  Thorpe  and  Rodger  having  shown  that  viscosities  depend 
largely  upon  the  frictional  surfaces  of  the  physical  particles  in  any 
solution,  it  followed  that  if  these  surfaces  were  increased  or  diminished 
by  any  means  whatsoever,  there  would  result  a  corresponding  increase 
or  decrease  in  the  fluidity  of  the  medium.  The  work  of  Jones  and 
Murray,  as  already  stated,  had  brought  out  the  fact  that  on  mixing 
two  highly  associated  liquids  a  material  diminution  in  the  association 
of  both  takes  place.  This  would  result  in  an  increase  in  the  number  of 
ultimate  particles  in  a  given  volume,  with  a  corresponding  decrease  in 
their  size,  and  therefore  an  increase  in  their  frictional  surfaces  which 
would  be  exposed  to  one  another.  Consequently,  the  viscosity  curves 
for  the  various  mixtures  should  pass  through  a  maximum,  the  position 
of  which  would  depend  upon  the  relative  effects  of  the  solvents  upon 
one  another. 

Quite  recently  Jones  and  Davis,  studying  viscosity,  tested  the  results 
obtained  by  Jones  and  Murray,  using  the  same  mixtures  which  they 
had  employed — acetic  acid  and  water,  formic  acid  and  water,  and  formic 
acid  and  acetic  acid.  They  obtained  results,  recorded  later  in  this 
chapter,  which  were  in  perfect  accord  with  the  earlier  work  of  Jones 
and  Murray.  It  was  shown  that  the  curve  for  the  viscosity  of  acetic 
acid  and  water  passes  through  a  well-defined  maximum  towards  the 
acetic  acid  end  of  the  curve,  showing  the  large  effect  of  water  on  the 
association  of  acetic  acid;  while  the  curve  for  formic  acid  and  acetic 


Viscosities  of  Caesium  Salts.  5 

acid  exhibited  only  a  slight  maximum  near  its  center,  which  showed 
that  the  two  liquids  possessed  about  the  same  small  effect,  each  on  the 
association  of  the  other.  On  the  other  hand,  the  curve  for  formic  acid 
and  water  was  practically  a  straight  line,  indicating  that  neither  solvent 
altered  appreciably  the  association  of  the  other. 

The  reason  for  the  maximum  in  such  viscosity  curves  and  their  rela- 
tion to  conductivity  having  been  explained,  the  effect  of  certain  salts 
in  lowering  the  viscosity  of  the  solvent  remained  to  be  interpreted. 
Potassium,  rubidium,  and  caesium  salts,  as  has  already  been  stated, 
were  known  to  lower  the  viscosity  of  water.  These  elements  occupy 
the  highest  maxima  on  the  atomic  volume  curve  and,  as  Wagner  has 
shown,  they  possess  negative  viscosity  coefficients  in  water  which  vary 
directly  as  their  volumes.  The  explanation  offered  by  Jones  and 
Veazey  to  account  for  these  facts  was  that  salts  of  these  metals  lower 
the  viscosity  of  the  solvent  by  introducing  into  it  ions  which  are  so 
large  that  when  mixed  with  the  molecules  of  the  solvent  they  lower  the 
total  frictional  surfaces  exposed  to  one  another.  That  certain  salts  of 
potassium  do  not  produce  this  effect  is  due  to  the  fact  that  the  viscosity 
of  the  solution  is  also  a  function  of  the  anions  as  well  as  of  the  cations. 
When  the  volume  of  the  anion  is  very  small,  the  negative  effect  of  the 
cation  on  the  viscosity  of  the  solution  is  more  than  overcome,  and 
positive  viscosity  results. 

Applying  this  hypothesis  of  Jones  and  Veazey  to  mixed  solvents, 
Jones  and  his  co-workers  have  been  able  to  explain  a  number  of  facts 
which  otherwise  appear  to  be  inexplicable.  For  example,  it  has  been 
found  that  rubidium  halides  and  potassium  iodide  lower  the  viscosity 
of  binary  mixtures  of  the  alcohols  and  of  acetone  with  water  which 
contain  as  much  as  50  per  cent  water ;  while  the  corresponding  increase 
in  viscosity  takes  place  in  all  mixtures  containing  the  larger  percentage 
of  alcohol.  The  explanation  of  this  fact  follows  at  once  from  what 
has  been  stated  above.  These  salts  lower  the  viscosity  of  water  and 
increase  the  viscosity  of  the  other  solvents  in  which  they  dissolve.  On 
mixing  the  two,  as  soon  as  the  association  of  the  alcohol  is  sufficiently 
diminished  so  that  the  solvent  aggregates  become  smaller  than  the  par- 
ticles of  the  dissolved  substances,  the  total  frictional  surface  between 
the  two  becomes  less  and  negative  viscosity  results. 

In  glycerol  as  a  solvent,  not  only  were  rubidium  salts  and  some 
potassium  salts  found  to  lower  the  viscosity  of  the  solvent,  but  also  cer- 
tain of  the  ammonium  salts.  So  great  was  the  negative  viscosity  effect 
produced  by  concentrated  solutions  of  rubidium  salts,  that  an  increase 
in  the  conductivity  of  these  solutions  over  that  of  the  more  dilute 
solutions  was  noted,  although  the  ionization  of  the  more  dilute  solu- 
tions was  undoubtedly  greater  than  that  of  the  more  concentrated. 
Ammonium  salts  were  found  to  be  more  closely  allied  to  salts  of  rubi- 
dium than  to  those  of  potassium,  in  respect  to  their  effect  on  viscosity. 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


The  explanation  of  the  above  phenomena  observed  by  Jones  and  his 
co-workers  follows  directly  from  and  substantiates  the  suggestion  made 
by  Jones  and  Veazey. 

APPARATUS  AND  PROCEDURE. 

In  carrying  out  this  investigation,  the  same  viscosity  apparatus  was 
used  as  in  our  work  with  formamid  as  a  solvent,  recorded  in  Chapter  II 
of  this  monograph.  The  scheme  of  temperature  regulation  and  the 
methods  of  obtaining  the  results  were  also  identical  with  those  used  in 
that  work. 

The  caesium  salts  were  all  carefully  recrystallized  and  dried  at  135° 
before  using. 

The  mixed  solvents  were  made  up  by  volume  at  20°,  from  liquids 
purified  by  the  usual  methods,  and  all  solutions  in  these  liquids  were 
prepared  at  the  same  temperature. 

TABLE  1. — Viscosity  and  Fluidity  of  Caesium  Salts  in  Mixtures  of  Water  with  Methyl  Alcohol, 
Ethyl  Alcohol,  and  Acetone  at  15°,  25°,  and  S5°. 


Viscosities 

Fluidities 

. 

Mol. 

cone. 

1,15° 

1,25° 

17  35° 

V15" 

«.25° 

«.35° 

Caesium  Chloride  in  75  p.  ct. 
Methyl  Alcohol   with 

fO.50 
J0.25 
10.10 

0.01674 
0.01637 
0.01613 

0.01309 
0.01275 
0.01243 

0.01049 
0.01021 
0.00997 

59.74 
61.09 
62.00 

76.39 
78.43 
80.45 

95.33 
97.94 
100.30 

Water. 

[Solv. 

0.01594 

0.01235 

0.01032 

62.75 

80.97 

96.90 

Cnssium  Chloride  in  50  p.  ct. 
Methyl   Alcohol   with 
Water. 

fO.50 
JO.  25 

m.io 
|Sotr. 

0.02119 
0.02084 
0.02086 
0.02100 

0.01547 
0.01553 
0.01531 
0.01535 

0.01200 
0.01197 
0.01179 
0.01169 

47.19 
47.98 
47.94 
47.62 

64.64 
64.39 
65.32 
65.15 

83.33 
83.54 
84.82 
85.54 

Ccesium  Cfdoride  in  25  p.  ct. 
Methyl   Alcohol    with 

fO.50 
J0.25 
10.10 

0.01778 
0.01828 
0.01862 

0.01338 
0.01360 
0.01358 

0.01048 
0.01047 
0.01044 

59.74 
61.09 
62.00 

76.39 

78.43 
80.45 

95.33 
97.  M 

100.34 

|.Solv. 

0.01871 

0.01359 

0.01032 

62.74 

80.97 

102.23 

r  2 

0.01098 

0.00877 

0.0713 

91.07 

114.03 

140.25 

Ccesium  Chloride  in  Water  .  . 

4 
1    10 

0.01111 
0.01133 

0.00889 
0.00884 

0.0716 
0.0719 

90.01 
88.18 

112.48 
113.12 

139  .  66 
139.08 

[Solv. 

0.01134 

0.00891 

0.0721 

88.18 

112.23 

138.89 

f     2 

Ccesium  Nitrate  in  75  p.  ct. 

* 

Methyl   Alcohol   with 

1    10 

0  01244 

0  00987 

80  39 

101  27 

Water. 

[Solv. 

0.01594 

0.01235 

0.00978 

62.74 

80.97 

102.23 

\     2 

Caesium  Nitrate  in  50  p.  ct. 

4 

0  01520 

0  01177 

65  83 

84.96 

Methyl   Alcohol   with 
Water  . 

\    10 
[Solv. 

0.02072 
0.02100 

0.01543 
0.01535 

0.01182 
0.01169 

48.26 
47.62 

64.81 
65.15 

84.60 
85.54 

f     2 

0  01300 

0  01018 

76  92 

98  23 

Ccesium  Nitrate  in  25  p.  ct. 
Methyl   Alcohol   with 
Water 

j     4 
10 

0.01795 
0.01832 

0.01331 
0.01346 

0.01028 
0.01031 

55.71 
54.59 

75.13 
74.29 

97.28 
96.99 

[Solv. 

0.01871 

0.01359 

0.01032 

53.45 

73.45 

96.90 

Viscosities  of  Caesium  Salts. 


TAKI.I:  1 .—  \'ix<-n.<<ity  and  Fluidity  of  Caesium  Salts  in  Mixtures  of  Water  with  Methyl  Alcohol, 
Ethyl  Alcohol,  and  Acetone  at  15°,  25°,  and  35° — Continued. 


Viscosities. 

Fluidities 

. 

Mol. 

cone. 

>;150 

i)25° 

ij35° 

»>15° 

^25° 

<p35° 

'     2 

0.01075 

0.00858 

0.00703 

93.02 

116.55 

142.25 

<'ii'*nint  XifrtUc  in  Water.  .  . 

4 
10 

0.01102 
0.01122 

0.00871 
0.00885 

0.00705 
0.00720 

90.74 
89.13 

114.81 
112.99 

141.84 
138.89 

.Solv. 

0.01184 

0.00891 

0.00720 

88.18 

112.23 

138.89 

Caesium  Chloride  in  75  p.  ct. 
Ethyl     Alcohol     with 

'     2 
4 
10 

0.02777 
0.02783 
0.02761 

0.02061 
0.02039 
0.01994 

0.01590 
0.01561 
0.01500 

36.01 
35.93 
36.22 

48.52 
49.04 
50.15 

62.89 
64.06 
66.67 

\\  :iter  . 

[Solv. 

0.02762 

0.01997 

0.01537 

35.21 

50.08 

65.06 

Ctrsium  Chloride  in  50  p.  ct. 
Ethyl     Alcohol     with 

'     2 
4 
10 

0.03176 
0.03346 
0.03359 

0.02245 
0  .  02283 
0.02274 

0.01635 
0.01666 
0.01648 

31.49 
29.89 
29.77 

44.54 
43.80 
43.98 

61.16 
60.02 
60.68 

Water  . 

.Solv. 

0  .  03400 

0  .  02286 

0.01618 

29.41 

43.73 

61.80 

(  VNI'MBI  Chluride  ill  25  p.  ct. 
Ethyl     Alcohol     with 

'     2 
4 
10 

0.02381 
0.02472 
0.02522 

0.01694 
0.01738 
0.01736 

0.01262 
0.01282 
0.01276 

42.00 
40.45 
39.65 

59.03 
57.54 
57.60 

79.24 
78.00 
78.37 

\\  ater  . 

[Solv. 

0.02585 

0.01760 

0.01270 

38.68 

56.82 

78.74 

f     2 

0.01098 

0.0877 

0.0713 

91.07 

114.03 

140.25 

Ccesium  Chloride  in  Water  .  . 

1      4 
1    10 

0.01111 
0.01133 

0.0889 
0.0884 

0.0716 
0.0719 

90.01 
88.18 

112.48 
113.12 

139.66 
139.08 

ISolv. 

0.01134 

0.0891 

0.0720 

88.18 

112.23 

138.89 

Caesium  Nitrate  in  75  p.  ct. 

f     2 
4 

Tnrol 

Ethyl     Alcohol     with 

i  n 

HloOl 

'     1  1 

Water  . 

10 

[Solv. 

uoie. 

'  — 

C     2 

('cesium  Nitrate  in  50  p.  ct. 

4 

Ethyl     Alcohol     with 

1    10 

Water  . 

[Solv. 

f     2 

Ctesium  Nitrate  in  25  p.  ct. 
Ethyl     Alcohol     with 

4 
1    10 

0.02424 
0.02526 

0.01697 
0.01745 

0.01261 
0.01276 

41.25 
39.59 

58.93 
57.31 

79.37 

78.37 

Water  . 

(Solv. 

0.02585 

0.01760 

0.01270 

38.68 

56.82 

78.74 

f     2 

0.01075 

0.00858 

0.00703 

93.02 

116.55 

142.25 

Cceiium  Nitrate  in  Water  .  .  . 

4 
1    10 

0.01102 
0.01122 

0.00871 
0.00885 

0.00705 
0.00720 

90.74 
89.13 

114.81 
112.99 

141.84 
138.89 

(Solv. 

0.01134 

0.00891 

0.00720 

88.18 

112.23 

138.89 

f     2 

Ccesium  Chloride  in  75  p.  ct. 

4 

0.01188 

0.00930 

0.00764 

84.18 

107  .  52 

130.89 

Acetone  with  Water. 

10 

0.01135 

0.00904 

0.00740 

88.11 

110.56 

135.08 

[Solv. 

0.01125 

0.00896 

0.00732 

88.88 

111.66 

136.59 

f     2 

0.01785 

0.01354 

0.01060 

56.02 

73.86 

94.34 

Ccesium  Chloride  in  50  p.  ct. 

4 

0.01783 

0.01337 

0.01038 

56.09 

74.79 

96.34 

Acetone  with  Water  . 

10 

0.01777 

0.01316 

0.01029 

56.27 

76.05 

97.18 

(.Solv. 

0.01766 

0.01306 

0.01009 

56.63 

76.57 

99.11 

f     2 

0.01647 

0.01250 

0.00983 

60.72 

80.00 

101.71 

Caesium  Chloride  in  25  p.  ct. 

4 

0.01667 

0.01254 

0.00979 

59.99 

79.74 

102.10 

Acetone  with  Water. 

10 

0.01677 

0.01246 

0.00978 

59.63 

80.26 

102.26 

[Solv. 

0.01690 

0.01248 

0  .  00964 

59.17 

80.12 

103  .  78 

Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


TABLE  1. — -Viscosity  and  Fluidity  of  Caesium  Salts  in  Mixtures  of  Water  with  Methyl  Alcohol, 
Ethyl  Alcohol,  and  Acetone  at  15°,  20°,  and  35°— Continued. 


Mol. 

Viscosities 

Fluidities 

I. 

cone. 

,15° 

7,25° 

J735" 

v>15° 

*>25° 

*>35° 

[     2 

0.01098 

0.00877 

0.0713 

91.07 

114.03 

140.25 

Caesium  Chloride  in  Water  .  . 

4 
10 

0.01111 
0.01133 

0.00889 
0.00884 

0.0716 
0.0719 

90.01 
88.18 

112.48 
113.12 

139.66 
139.08 

[Solv. 

0.01134 

0.00891 

0.0720 

88.18 

112.23 

138.89 

f     2 

CcBsium  N  itrate  in  75  p.  ct. 

4 

Acetone  with  Water. 

1    10 

0.01143 

0.00904 

0.00737 

87.48 

110.62 

135.76 

[Solv. 

0.01119 

0.00878 

0.00680 

89.37 

113.90 

147.12 

f     2 

Caesium  Nitrate  in  50  p.  ct. 

4 

0.01759 

0.01320 

0.01030 

56.85 

75.76 

97.09 

Acetone  with  Water. 

1    10 

0.01756 

0.01307 

0.01017 

56.95 

76.45 

98.33 

[Solv. 

0.01767 

0.01304 

0.01007 

56.59 

76.69 

99.30 

f     2 

0.01610 

0.01216 

0.00965 

62.11 

82.84 

103.63 

Ccesium  Nitrate  in  25  p.  ct. 

4 

0.01655 

0.01229 

0.00971 

60.42 

81.37 

103.02 

Acetone  with  Water  . 

1    10 

0.01665 

0.01244 

0.00968 

60.06 

SO.  39 

103.35 

[Solv. 

0.01673 

0.01232 

0.00962 

59.77 

81.18 

102.00 

\     2 

0.01075 

0.00858 

0.00703 

93.02 

116.55 

142.25 

Caesium  Nitrate  in  Water  .  .  . 

4 
10 

0.01102 
0.01122 

0.00871 
0.00885 

0.00705 
0.00720 

90.74 
89.13 

114.81 
112.99 

141.84 
138.89 

[Solv. 

0.01134 

0.00891 

0.00720 

88.18 

112.23 

138.89 

DISCUSSION  OF  RESULTS. 

As  can  be  seen  from  table  1,  measurements  have  been  made  of 
the  viscosities  of  solutions  of  these  salts  at  15°,  25°,  35°  in  mixtures 
of  25,  50,  and  75  per  cent  of  each  of  the  three  solvents  with  water. 
Caesium  salts,  with  the  possible  exception  of  the  carbonate,  are  prac- 
tically insoluble  in  the  alcohols  and  in  acetone;  and  in  a  number  of 
cases  it  was  impossible  to  obtain  solutions  more  concentrated  than  one- 
fourth  normal  in  the  75  per  cent  mixtures.  In  fact,  caesium  nitrate  was 
not  more  soluble  than  tenth-normal  in  75  per  cent  methyl  alcohol  and 
75  per  cent  acetone,  while  it  is  impossible  to  obtain  even  that  con- 
centration in  75  per  cent  ethyl  alcohol.  One  peculiarity  previously 
noted  in  the  case  of  rubidium  chloride  in  75  per  cent  acetone  was  also 
observed  with  caesium  chloride.  A  one-half  normal  concentration  of 
this  salt  could  easily  be  prepared,  but  with  a  partial  separation  of  the 
acetone  from  the  solvent.  However,  on  cooling  below  10°  a  perfectly 
homogeneous  mixture  again  resulted,  which  separated  again  into  two 
layers  on  warming.  The  nitrate  at  this  concentration  merely  remained 
partially  undissolved.  The  results  obtained  are  tabulated  for  each 
salt  under  the  heads  of  the  different  solvent  mixtures,  the  blank 
spaces  indicting  that  the  salt  was  insoluble  at  the  concentration  in 
question.  In  general,  both  of  the  salts  lower  the  viscosity  of  water  and 
of  the  25  per  cent  mixture  of  all  the  solvents,  and  increase  the  viscosity 


Viscosities  of  Caesium  Salts.  9 

of  the  75  per  cent  mixture.  These  facts  are  in  perfect  accord  with  the 
results  obtained  in  earlier  work  with  salts  of  potassium  and  rubidium. 
In  the  50  per  cent  mixtures  the  two  salts  manifest  somewhat  different 
behavior.  The  chloride  in  50  per  cent  methyl  alcohol  shows  an 
increase  in  viscosity  at  all  temperatures  for  the  most  dilute  solutions, 
and  a  tendency  of  the  more  concentrated  solutions  to  pass  over  into 
negative  viscosity.  The  nitrate  in  the  same  solvent  showed  a  decided 
negative  viscosity  effect. 

In  50  per  cent  ethyl  alcohol  a  decrease  in  the  viscosity  of  the  solu- 
tions as  compared  with  that  of  the  solvent  was  noted  at  the  lower 
temperature,  a  transition  to  positive  viscosity  taking  place  at  higher 
temperatures.  Again,  the  nitrate  decreases  the  viscosity  of  the  solvent 
at  all  temperatures. 

With  regard  to  the  50  per  cent  acetone-water  mixture,  an  increase 
in  viscosity  at  all  temperatures  is  to  be  noted  in  the  case  of  the  chloride, 
while  the  nitrate  in  50  per  cent  shows  a  tendency  to  decrease  the 
viscosity  of  the  solvent,  although  it  increases  it  at  the  higher  tempera- 
tures. The  difference  in  the  effect  produced  by  these  salts,  in  com- 
parison with  that  produced  by  the  salts  of  rubidium,  lies  in  the  shifting 
of  the  transition-point  from  negative  to  positive  viscosity  towards  the 
mixtures  containing  the  larger  percentage  of  acetone.  This  follows 
from  the  different  molecular  volumes  of  the  cations  of  rubidium  and 
csesium.  When  acetone  and  water  are  mixed  the  principal  changes 
take  place  in  the  association  of  the  acetone.  Consequently,  until 
considerable  water  has  been  introduced  the  solvent  particles  are  quite 
large,  so  that  even  csesium  nitrate  having  the  largest  negative  viscosity 
coefficient  in  water  increases  the  viscosity  of  the  75  per  cent  mixture  and 
even  of  the  50  per  cent  mixture,  except  at  lower  temperatures.  For 
rubidium  salts,  no  examples  of  negative  viscosity  are  to  be  noted  until 
the  37.5  per  cent  mixture  is  reached.  No  data  are  available  for  com- 
parison with  rubidium  salts  in  alcohol-water  mixtures,  but  it  has  been 
shown  that  potassium  iodide  increases  the  viscosity  of  these  mixtures  to 
that  concentration  containing  40  per  cent  alcohol.  But  in  this  instance, 
csesium  nitrate  decreases  the  viscosity  of  the  50  per  cent  mixture,  posi- 
tive viscosity  manifesting  itself  only  in  the  75  per  cent  mixture. 

This  shifting  of  the  transition-point  from  negative  to  positive  vis- 
cosity towards  the  more  concentrated  solvents  (regarding  water  as  the 
diluent),  with  increase  in  the  molecular  volume  of  the  salt,  brings  out 
clearly  the  gradual  breaking-down  of  the  associated  molecules  into- 
smaller  particles  with  greater  frictional  surfaces;  the  difference  in  the 
sizes  of  the  particles  at  different  points  on  the  dilution  curve  is  clearly 
seen  from  the  effect  produced  by  salts  of  different  molecular  volumes. 

The  apparent  transition  from  negative  to  positive  viscosity  with  rise 
in  temperature,  noted  in  certain  instances,  would  seem  to  indicate 
either  a  polymerization  of  the  salt  or  else  a  solvent  envelope  which  is 


10  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

broken  up  with  rise  in  temperature.     Since  these  salts  are  the  least 
solvated,  the  first  assumption  is  apparently  the  more  plausible. 

Meaurements  are  now  in  progress  of  the  effect  of  these  salts  on  the 
viscosity  of  glycerol  and  of  glycerol-water  mixtures. 

THE  VISCOSITIES  OF  BINARY  MIXTURES  OF  THE  ASSOCIATED 
LIQUIDS,  WATER,  FORMIC  ACID,  AND  ACETIC  ACID. 

Jones  and  Murray1  showed  in  1903  that  when  two  associated  liquids 
are  mixed  each  diminishes  the  association  of  the  other.  They  deter- 
mined the  molecular  weight  of  water  in  formic  acid  on  the  one  hand 
and  in  acetic  acid  on  the  other.  Also  the  molecular  weight  of  formic 
acid  in  water,  and  in  acetic  acid,  and  finally  the  molecular  weight  of 
acetic  acid  in  water  and  in  formic  acid. 

They  found  that  the  molecular  weight  of  water  in  formic  acid  varied 
from  19.7  at  dilution  0.93  N  to  21.9  at  6.18  N,  showing  that  the  formic 
acid  diminished  slightly  the  association  of  the  water. 

The  molecular  weight  of  water  in  acetic  acid  varied  from  21.7  at  0.64 
N  to  38.8  at  12.65  N.  This  showed  that  the  acetic  acid  diminished 
greatly  the  association  of  the  water. 

The  molecular  weight  of  acetic  acid  in  water  varied  from  55.4  at  0.17 
N  to  72.1  at  7.06  N,  showing  that  water  diminished  very  appreciably 
the  association  of  acetic  acid. 

The  molecular  weight  of  acetic  acid  in  formic  acid  varied  from  61.9 
at  0.18  N  to  83.8  at  9.17  N.  The  association  of  the  acetic  acid  was  thus 
diminished  considerably  by  the  formic  acid. 

The  molecular  weight  of  formic  acid  in  water  varied  from  45.2  at  0.38 
N  to  51.0  at  6.16  N,  showing  that  water  had  very  little  effect  on  the 
association  of  the  formic  acid;  and,  finally,  the  molecular  weight  of 
formic  acid  in  acetic  acid  varied  from  50.4  at  0.82  N  to  65.7  at  8.26  N. 

If  we  take  into  account  the  dissociating  action  of  the  water  on 
formic  and  acetic  acids,  we  are  justified  in  concluding  that  water  and 
formic  acid  affect  each  other's  association  very  little;  acetic  acid  and 
formic  acid  affect  each  other's  association  considerably,  while  both 
water  and  acetic  acid  have  a  marked  effect  each  on  the  association 
of  the  other.  These  facts  were  used  by  Jones  and  Veazey2  to  explain  the 
increase  in  viscosity  which  takes  place  when  water  and  alcohol  are  mixed. 
These  are  both  strongly  associated  solvents,  and  each,  therefore,  very 
probably  diminishes  appreciably  the  association  of  the  other.  From 
a  smaller  number  of  larger  molecules  of  each  solvent,  we  have  a  larger 
number  of  smaller  molecules  of  each  solvent  produced.  The  surfaces 
of  the  molecules  present  are  therefore  increased,  and,  consequently, 
the  friction  of  these  molecules  as  they  move  over  one  another.  Vis- 
cosity is  a  function  of  the  surface  friction  of  the  molecules. 

'Amer.  Chem.  Journ.,  30,  193  (1903). 

•Zeit.  phys.  Chem.,  61,  641;  62,  44  (1908);  Carnegie  Inst.  Wash.  Pub.  No.  80,  170  (1907). 


Viscosities  of  Caesium  Salts. 


11 


In  the  same  manner  Jones  and  Veazey  were  able  to  explain  why  it 
is  that  salts  of  potassium,  rubidium,  and  caesium  lower  the  viscosity 
of  water  and  other  solvents1  in  which  they  are  dissolved.  The  large 
atomic  volumes  of  these  cations,  when  mixed  with  the  molecules  of 
water,  diminish  the  frictional  surfaces  which  come  in  contact,  and, 
consequently,  diminish  the  viscosity. 

Having  found  the  above  action  of  water,  formic  acid,  and  acetic  acid, 
each  on  the  association  of  the  other,  the  question  arose,  What  would  be 
the  effect  of  each  on  the  viscosity  of  the  other?  If  the  above  explana- 
tions offered  by  Jones  and  Veazey  were  correct,  then  two  liquids,  like 
water  and  acetic  acid,  which  diminished  each  other's  association,  ought 
to  increase  the  viscosity  of  one  another — the  viscosity  of  the  mixture 
should  be  greater  than  that  of  either  pure  liquid  separately. 

The  viscosities  of  binary  mixtures  of  the  above-named  three  liquids 
were  measured  at  15°  and  25°.  Water  was  regarded  as  the  solvent  for 
formic  acid  and  acetic  acid,  and  solutions  of  these  acids  in  water  were 
prepared  containing  by  volume  10,  20,  30,  40,  50,  60,  70,  80  and  90  per 
cent.  The  acids  themselves  contained  somewhat  less  than  1  per  cent 
of  water.  The  viscosity  and  the  fluidity  (reciprocal  of  viscosity)  data 
are  given  in  tables  2  to  4. 

TABLE  2. — Viscosities  and  Fluidities  of  Formic  Add  in  Water  at  16°  and  25°. 


Per  cent  HCOOH. 

17  15° 

«,15° 

i)25° 

«,25° 

(H,O) 

0  01134 

88  18 

0  00891 

112  23 

10  

0  01215 

82  31 

0  00932 

107  30 

20 

0  01282 

78  00 

0  01014 

98  62 

30  

0  01339 

74  68 

0  01072 

93  28 

40 

0  01408 

71  02 

0  01135 

88  11 

50  

0  01469 

68  07 

0  01202 

83  20 

60  . 

0  01591 

62  85 

0  01287 

77  70 

70  

0.01693 

59  07 

0  01371 

72  94 

80  . 

0  01803 

55  46 

0  01452 

68  87 

90 

0  01914 

52  25 

0  01546 

64  68 

(HCOOH) 

0  01963 

50  94 

0  01571 

63  65 

The  viscosity  data  at  15°  are  plotted  in  curves,  fig.  1,  and  the  vis- 
cosity data  at  25°  in  curves,  fig.  2.  The  viscosities  of  acetic  acid  in 
water  pass  through  a  well-defined  maximum,  which,  before  we  carried 
out  a  single  measurement  of  the  viscosities  of  mixtures  of  these  two 
solvents,2  we  predicted  would  be  the  case  from  the  molecular-weight 
determinations  of  this  acid  in  water  made  by  Jones  and  Murray. 

The  viscosities  of  formic  acid  in  acetic  acid  pass  through  a  slight 
maximum,  as  would  be  expected  from  the  effect  of  each  of  these  sol- 
vents on  the  molecular  weight  of  the  other. 

'Jones  and  Davis:  Zeit.  phys.  Chem.,  81,  68  (1912) ;  Carn.  Inst.  Wash.  Pub.  No.  180, 179  (1913). 

"Since  completing  our  work  we  find  that  a  few  measurements  of  the  viscosities  of  mixtures  of 
acetic  acid  in  water  had  been  made  by  Dunstan  and  Thole:  Journ.  Chem.  Soc.,  85,  825  (1904); 
95,  1560  (1809). 


12  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

TABLE  3. — Viscosities  and  Fluidities  of  Acetic  Add  in  Water  at  15"  and  85°. 


Per  cent  CHjCOOH. 

i)15" 

tp  15° 

ij25° 

*>25° 

(H2O)'  

0  01134 

88  18 

0  00891 

112  23 

10  

0  01368 

73  10 

0  01059 

94  43 

20.  .  . 

0  01626 

61  50 

0  01244 

30  

0  01897 

5°  72 

0  01446 

40  

0  02143 

46  66 

0  01624 

61  58 

50. 

0  02416 

41  39 

0  01818 

60  

0  02682 

37  29 

0  0°015 

70  

0  02935 

34  07 

0  0'7219 

45  07 

80  

0  03068 

32  60 

0  02318 

43  14 

85  

0  03033 

32  90 

0  0^92 

90  

0  02786 

35  89 

0  0^115 

95  

0  02243 

44  58 

0  01775 

56  34 

CHtCOOH  

0  01410 

70  92 

0  01174 

85  18 

'Values  for  water  taken  from  Thorpe  and  Rodger:  Phil.  Trans.,  185A,  449  (1894) 


TABLE  4. — Viscosities  and  Fluidities  of  Formic  Acid  in  Acetic  Acid  at  13°  and  S5°. 


Per  cent  HCOOH. 

»;150 

»,150 

i/25° 

P25° 

(CH,COOH)  

0  01410 

70  92 

0  01174 

CC   10 

10  

0  01558 

64  19 

0  01286 

20  

0  01701 

58  79 

0  01391 

30  

0  01792 

55  80 

0  01463 

40  

0  01883 

53  11 

n  m  <iftfi 

50  

0  01942 

->]  .(<) 

0  01564 

60  

0  01984 

50  40 

0  01587 

70  

0  02012 

49  70 

Onifin7 

80  

0  02002 

49  95 

0  01607 

90  

0  01967 

50  84 

0  01582 

(HCOOH)  

0  01963 

50  94 

0  01571 

300 


250 


O 
X 


'S200 

8 

i 


ISO 


100 


.5 


^ 


\nai 


10     20     30    4O     50     60    70     80    9O     1OO 

Concentration 
Fio.  1. 


E50 

200 
O 

- 

5 

& 

- 

25° 

*$' 

' 

\ 

fy 

* 

5Ld-'! 

,a-c-e< 

15 

X 

£•'50 

1 

100 

50 

$ 

fe\ 

I--- 

? 

,< 

r-< 

' 

|4»>C 

r  ^( 

c 

1 

A  t« 

<•«*$ 

ffi 

>-' 

-A-\ 

\"\ 

10     20     30     4O     SO    6O     TO      80    90     ICX 

Concentration 

Fio.  2. 


Viscosities  of  C cesium  Salts.  13 

The  viscosities  of  mixtures  of  water  and  formic  acid,  on  the  other 
hand,  fall  almost  on  a  straight  line.  This,  again,  is  just  what  would  be 
expected,  from  the  fact  that  neither  solvent  alters  appreciably  the 
molecular  weight  of  the  other. 

These  results  are,  then,  all  in  keeping  with  the  suggestion  made  by 
Jones  and  Veazey,  that  the  increase  in  viscosity  which  results  when 
associated  liquids  which  diminish  each  other's  association  are  mixed,  is 
due  to  the  larger  number  of  smaller  parts  that  are  present. 

These  results  are  in  perfect  accord  with  their  suggestion  as  to  the 
cause  of  the  diminution  in  the  viscosity  of  water  produced  by  salts 
whose  cations  have  very  large  ionic  volumes,  such  as  salts  of  potassium, 
rubidium,  and  caesium. 

SOME  NEW  FORMS  OF  APPARATUS. 

A  SUBSTITUTE  FOR  THE  TWIN-BULB  TRAP  IN  TOLUENE-MERCURY 
THERMO-REGULATORS. 

Toluene,  on  account  of  its  high  coefficient  of  expansion,  is  to  be  pre- 
ferred to  all  other  liquids  for  use  in  thermo-regulators.  Since,  however, 
it  is  practically  a  nonconductor  and  quite  volatile,  it  is  ordinarily  used 
with  mercury  for  the  contact  in  all  electrically  operated  thermostats. 

In  order  to  use  these  two  liquids  together,  the  common  form  of  appa- 
ratus hitherto  employed  has  been  the  twin-bulb  device,  which,  however, 
has  the  following  disadvantages: 

1.  It  is  very  fragile  and  can  only  be  made  by  an  expert  glass-blower. 

2.  When  the  mercury  level  in  the  capillary  has  been  once  adjusted 
for  any  given  temperature  and  the  toluene  reservoir  sealed,  the  regu- 
lator is  practically  useless  for  any  higher  temperatures  without  opening 
the  reservoir  and  removing  the  excess  of  toluene  in  order  to  preserve 
the  equilibrium  in  the  two  bulbs. 

3.  It  is  difficult  to  prevent  the  toluene  from  finally  creeping  around 
between  the  mercury  and  the  glass  walls  into  the  capillary  and  fouling 
the  contact  surface  of  the  mercury,  since  the  same  continuous  tube 
contains  both  the  toluene  and  the  mercury. 

To  overcome  these  difficulties  as  far  as  possible,  the  apparatus 
illustrated  in  figure  3  has  been  devised.  This  consists  of  a  bulb  a 
attached  at  the  bottom  by  the  tube  b  to  the  toluene  reservoir,  which 
may  be  of  any  desired  form.  Exactly  opposite  to  b  is  a  corresponding 
tube  (c),  which  carries  the  capillary  and  sealed-in  platinum  contact. 

The  interior  construction  of  a  is  as  shown  in  the  figure.  The  small 
tubes  e  and/ are  prolongations  of  c  and  b,  respectively,  having  a  length 
nearly  equal  to  the  diameter  of  the  bulb.  These  tubes  are  inclined  at 
an  angle  of  about  30°  to  the  perpendicular.  The  short  side  tube  d  is 
used  in  filling  the  regulator  and  is  in  exact  alinement  with  /. 

To  prepare  the  regulator  for  use,  mercury  is  poured  in  through  the 
capillary  until  the  bulb  a  is  from  one-half  to  three-fourths  filled.  A 


14 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


tube  drawn  out  to  a  small  diameter  is  then  inserted  through  d  and  / 
into  b,  and  toluene  allowed  to  run  in  until  the  reservoir  is  filled  and 
also  the  space  in  a  above  the  mercury.  The  tube  d  is  then  sealed, 
and  by  adding  or  withdrawing  mercury  from  the  capillary  the  contact 
level  may  be  adjusted  for  any  desired  temperature. 

By  this  method  a  toluene-mercury  regulator  is  constructed  with  a 
range  of  50°  or  more,  all  adjustments  being  made  through  the  capillary, 
which  should  be  of  from  1.5  to  2  mm.  bore,  according  to  the  capacity 
of  the  toluene  reservoir. 

All  danger  of  the  toluene  creeping  into  the  capillary  is  removed,  since 
the  tubes  e  and/ terminate  in  the  same  liquid  which  is  contained  in  them. 

This  trap  is  compact,  not  easily  broken,  and  is  comparatively  simple 
to  construct.  In  addition  to  this,  it  lies  within  the  plane  of  the  reser- 
voir and  contact  tubes. 


FIG.  3.  FIG.  4. 

A  NEW  FORM  OF  PYKNOMETER  FOR  LIQUIDS.' 

The  investigators,  P.  B.  Davis  and  L.  S.  Pratt,  have  recently  had  occa- 
sion to  carry  out  a  large  number  of  density  determinations  with  various 
liquids.  For  this  purpose  a  modified  form  of  the  Sprengel  pyknometer, 
as  described  by  Jones  and  Bingham2  and  Jones  and  Veazey,3  was  first 
adopted,  but  was  found  to  have  the  following  disadvantages: 

1.  In  instruments  of  large  capacity  (10  to  20  c.c.)  the  long  capillary 
side-arm  adds  greatly  to  the  weight  and  hence  impairs  the  accuracy  of 
the  pyknometer. 

'Devised  with  L.  S.  Pratt.     'Araer.  Chem.  Journ.,  34, 48(1905).      'Zeit.phys.Chem.,61,641  (1908). 


Viscosities  of  Caesium  Salts.  15 

2.  The  fragility  of  the  older  form  makes  rapid  handling  practically 
impossible. 

3.  The  instrument  is  difficult  to  dry  and  polish  before  weighing. 
To  remove  these  defects  the  form  shown  in  fig.  4  was  devised,  and 

has  been  found  to  be  more  convenient  and  accurate  than  the  old  form. 
In  the  figure  the  symbols  have  the  following  significance : 
(a)  A  reservoir  of  thin-walled  glass  tubing  which  may  be  of  any 

desired  capacity,  those  in  use  in  this  laboratory  holding,  respectively, 

about  6,  10,  and  20  c.c. 

(&)  Tapering  tube  of  thin  glass  reaching  almost  to  the  bottom  of  (a), 

and  having  a  bore  at  the  open  end  equal  to  or  slightly  greater  than  that 

of  the  capillary  (0.5  to  0.75  mm.).     This  tube  is  of  use  in  drying  and 

filling  the  instrument. 

(c)  Outlet  tube  sealed  on  to  the  end  of  6.    This  is  drawn  out  slightly 
at  the  end  and  is  bent  at  an  angle  of  60°  close  to  the  reservoir. 

(d)  Slight  enlargement  in  c  which  holds  in  position  the  wire  used 
to  suspend  the  pyknometer  on  the  balance. 

(e)  Expansion  bulb  of  sufficient  size  to  accommodate  the  increased 
volume  of  any  liquid  when  warmed  10°.     If  the  instruments  are  used 
only  at  25°  this  bulb  need  be  only  very  small. 

(/)  Fine  line  etched  around  the  capillary  at  the  lower  limit  of  the 
taper  in  the  bulb  e. 

In  adjusting  the  instrument  when  filled,  a  rubber  tube  is  attached 
to  the  horizontal  capillary  above  e,  and  the  liquid  blown  out  gently 
until  the  level  in  e  just  reaches  the  mark  /.  The  excess  of  liquid 
is  then  removed  from  the  end  of  c  with  filter  paper.  On  removing 
the  pressure,  the  level  of  the  liquid  in  c  falls  until  it  comes  to  rest 
near  the  bend  in  the  capillary,  thereby  lessening  the  danger  from  evap- 
oration or  accidental  spilling  after  adjustment. 

The  advantages  of  this  form  of  pyknometer  can  readily  be  seen,  e.  g., 
the  projecting  arms  are  both  short  and  hence  not  easily  broken.  Also, 
the  net  weight  has  been  materially  lessened,  since  the  amount  of  neces- 
sary capillary  has  been  reduced  to  a  minimum.  The  instrument  is 
also  much  easier  to  clean  and  dry  before  weighing  than  the  older  form. 

This  pyknometer  may  be  provided  with  ground-on  caps  when  used 
with  alcohol  or  other  volatile  liquids,  to  prevent  possible  evaporation. 


CHAPTER  II. 

THE  CONDUCTIVITY  AND  VICOSITY  OF  SOLUTIONS  OF  CERTAIN 
SALTS  IN  FORMAMID  AS  A  SOLVENT. 


BY  P.  B.  DAVIS  AND  W.  S.  PUTNAM. 


The  work  of  Jones  and  his  collaborators  in  the  field  of  non-aqueous 
solvents  and  in  mixtures  of  these  solvents  with  one  another,  has  thus 
far  necessitated  a  comprehensive  study  of  both  the  conductivity  and 
viscosity  of  typical  salts  in  methyl  and  ethyl  alcohols,  in  acetone,  and 
in  glycerol — as  well  as  in  binary  and  ternary  mixtures  of  these  solvents 
with  one  another  and  with  water.  A  complete  review  of  all  this  work 
is  to  be  found  in  Publication  No.  210  of  the  Carnegie  Institution  of 
Washington,  Chapter  VII. 

Because  of  the  somewhat  limited  solubility  of  electrolytes  in  most 
organic  solvents,  the  scope  of  this  work  has  necessarily  been  somewhat 
limited  in  the  case  of  the  pure  liquids  themselves,  This  is  especially 
true  with  acetone  and  the  alcohols.  Glycerol,  however,  notwithstand- 
ing its  high  viscosity,  proved  to  be  a  remarkably  good  solvent. 

Among  the  few  remaining  liquids  suitable  for  such  work,  formamid, 
the  one  used  in  this  investigation,  is  the  most  important.  The  fact 
that  it  has  been  studied  so  little  in  the  past  must  be  attributed  to  the 
difficulties  encountered  in  obtaining  a  product  pure  enough  for  con- 
ductivity purposes. 

Formamid,  the  first  member  of  the  series  of  acid  amids  of  which  aceta- 
mid  is  the  best  known  representative,  was  first  prepared  by  A.  W.  von 
Hoffman1  in  1863,  by  heating  anhydrous  ethyl  formate  with  dry 
ammonia  in  a  sealed  tube  for  some  time,  and  then  distilling  the  resulting 
product  in  a  partial  vacuum.  A  number  of  other  methods  of  prepara- 
tion have  been  devised,  among  the  earlier  ones  being  that  of  Behrend,2 
who  heated  a  dry  mixture  of  ammonium  formate  and  urea  to  140°  and 
purified  the  oily  residue  by  fractional  distillation  in  vacua.  More 
recently  Phelps  and  Deming3  improved  Hoffman's  method  by  carrying 
out  the  decomposition  of  the  ethyl  formate  at  a  temperature  below 
that  at  which  ammonium  formate  is  formed,  thereby  obtaining  nearly 
the  theoretical  yield.  In  addition  to  the  above,  probably  the  most 
common  method  of  preparing  formamid  is  that  depending  on  the  disso- 
ciation of  anhydrous  ammonium  formate  by  heat  in  an  atmosphere  of 
dry  ammonia,  as  described  for  the  first  time  also  by  Hoffman.4  By 
this  method,  with  the  latest  improvements  due  to  Freer  and  Sherman,6 
pure  formic  acid  is  neutralized  with  dry  ammonia  and  the  resulting 

'Journ.  Chem.  Soc.,  16,72(1863).    2Lieb.  Ann.,  128, 335 (1863).    'Centralblatt  II,  1604(1907). 
*Ber.  d.  deutsch.  chem.  Gesell.,  15,  980  (1882).     'Amer.  Chem.  Journ.,  20,  223  (1898). 

16 


Conductivities  and  Viscosities  in  Formamid.  1 7 

formate  heated  for  several  hours  to  180°  in  an  atmosphere  of  ammonia. 
In  distilling  the  resulting  brownish  liquid  under  diminished  pressure, 
a  yield  of  80  per  cent  of  the  theoretical  was  obtained.  At  lower  pres- 
sures (0.5  to  2  mm.)  a  better  yield  resulted,  since  in  this  way  the  thermal 
decomposition  was  reduced  to  a  minimum. 

PREVIOUS  WORK  IN  FORMAMID. 

Until  quite  recently  no  investigation  of  the  physical-chemical  prop- 
erties of  this  remarkable  liquid  had  been  made.  A  brief  sketch  of  the 
recent  work  is  given  below. 

In  his  preliminary  paper  on  organic  solvents,  Walden  mentions  the 
marked  similarity  of  formamid  to  water  as  compared  with  other  liquids 
in  its  solvent  action  on  inorganic  salts,  and  in  later  papers  he  records 
some  of  its  physical  constants  in  comparison  with  other  amids  and  with 
water,  viz,  the  dielectric  constant  and  association  factors,1  specific 
conductivity  and  molecular  conductivity  of  the  normal  electrolyte 
N(C2H5)4I,2  and  the  viscosity.3  All  of  these  constants  were  obtained 
with  a  product  which,  as  will  be  shown  later,  was  far  from  pure,  although 
it  may  be  pointed  out  that  the  first  two  constants  are  not  appreciably 
affected  by  small  changes  in  the  purity  of  the  solvent. 

Turner  and  Merry,4  in  their  work  on  the  molecular  composition  of 
trivalent  nitrogen  compounds,  pointed  out  that  the  high  association 
factor  of  formamid  is  one  of  its  most  striking  characteristics,  and  they 
observed  its  similarity  to  water.  They  noted  also  that  the  association 
of  formamid  diminishes  more  rapidly  with  rise  in  temperature  than  that 
of  water,  and  suggested  that  the  solvent  power  for  salts  was  largely  due 
to  its  high  molecular  complexity. 

Somewhat  later,  Walden,5  in  a  study  of  the  temperature  coefficient 
K,  for  organic  liquids  in  the  Ramsay  and  Shields  equation  for  molecular 
surface-tension,  found  that  in  the  case  of  formamid  the  value  was  far 
below  that  for  non-associated  liquids  (2.12),  having  only  the  value 
0.594  to  0.710;  from  these  data  he  obtained  an  association  factor  in 
close  agreement  with  that  as  determined  by  Turner  and  Merry.4 

Dunstan  and  Kassel,6  while  determining  the  fluidity  of  various  binary 
mixtures,  measured  the  viscosity  of  mixtures  of  formamid  and  i-amyl 
alcohol  at  both  low  and  somewhat  elevated  temperatures,  and  noted  a 
pronounced  minimum  in  the  fluidity  curves  at  about  60  to  70  per  cent 
of  formamid  at  both  temperatures,  and  also  a  slight  maximum  at  about 
10  to  20  per  cent  formamid  for  the  lower  temperature. 

Rohler7  studied  the  solvent  properties  of  formamid  for  organic  salts 
and  also  the  electrolysis  of  its  solutions.  He  compared  the  density 
with  that  of  water  above  the  freezing-point  and  found  no  minimum, 

'Zeit.  phys.  Chem.,  46,  145,  175  (1906).    'Ibid.,  54  179(1905).    3Ibid.,  55,  230  (1906). 

Mourn.  Chem.  Soc.,  97,  2076  (1910). 

•Zeit. phys.  Chem. ,75, 555  (1910).   '/Wd.,76, 367  (1911).   'Zeit.  Elektrochem.,  16, 420  (1910). 


18  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

as  in  the  case  of  the  latter,  but  that  the  solvent  expanded  linearly  with 
rise  in  temperature,  the  variation  of  the  specific  volume  being  repre- 
sented by  V  =  0.8674  X  (10.000742<) . 

Rohler  also  found  that,  according  to  Faraday's  law,  copper  dissolves 
in  formamid  at  all  dilutions,  half  as  univalent  and  half  as  bivalent  ions, 
at  the  anode.  He  obtained  a  good  separation  of  copper  with  a  weak 
current,  but  with  a  stronger  current  the  metal  separated  as  a  dark 
slime.  Similar  results  were  obtained  with  lead,  zinc,  and  tin.  With 
all  of  these  metals  in  air,  more  metal  dissolved  at  the  anode  than 
separated  at  the  cathode.  The  electrolysis  of  nickel,  cobalt,  iron, 
aluminium,  and  magnesium  salts  yielded  no  appreciable  amount  of 
metal  at  the  cathode.  Rohler  also  noted  the  formation  of  metallic 
formamidates  and  of  crystalline  double  compounds,  which  will  be  con- 
sidered in  discussing  the  properties  of  formamid  as  compared  with 
those  of  water. 

In  a  quite  recent  paper  Walden1  gives  an  account  of  his  cryoscopic 
work  in  this  solvent.  In  his  preliminary  discussion  he  states,  as  his 
reason  for  adopting  the  freezing-point  rather  than  the  conductivity 
method  for  measuring  dissociations,  that  the  high  specific  conductivity 
of  the  solvent  (as  measured  by  him)  rendered  it  impossible  to  obtain 
accurate  data  by  the  latter  method.  However,  as  will  be  pointed  out 
later,  by  our  method  of  preparing  formamid  a  solvent  may  be  obtained 
with  a  specific  conductivity  comparable  with  that  of  water,  the  loss  of 
material  being  only  one-third  that  experienced  by  Walden. 

Walden  determined  the  freezing-point  constant  of  formamid,  and 
gives  the  value  35.0  as  the  mean  of  six  determinations,  using  urea, 
acetic  ester,  diethylsulphite,  ethyl  acetate,  mesityloxide,  and  nitrana- 
line,  respectively,  as  the  solute.  He  also  studied  the  dissociation  of  a 
number  of  electrolytes,  including  salts  and  both  strong  and  weak  acids, 
and  pointed  out  that  all  binary  salts  are  strongly  ionized  at  relatively 
high  concentrations,  the  ionization  increasing  slowly  to  the  limit  a  =  100, 
as  in  water;  the  limit,  however,  being  reached  at  a  smaller  dilution  in 
formamid.  The  same  was  found  to  hold  for  acids,  except  in  the  case 
of  those  where  combination  took  place  between  solvent  and  solute,  as 
in  the  case  of  some  of  the  strong  acids,  the  values  found  being  smaller 
than  those  for  water. 

FORMAMID  AS  A  SOLVENT. 

Formamid  is  a  clear,  colorless,  and  somewhat  viscous  liquid,  melting 
at  18°,  and  boiling  under  atmospheric  pressure  with  partial  decomposi- 
tion at  200°  to  212°.  It  reacts  neutral  to  litmus  and  is  quite  hydrc- 
scopic,  undergoing  a  slow  hydrolysis  at  ordinary  temperatures  into 
ammonium  formate.  This  solvent  is  the  most  closely  allied  to  water 
in  its  properties  of  all  the  organic  solvents.  The  two  are  miscible  in 

'Bull.  Imp.  Acad.  Sci.,  St.  Petersburg  (1911). 


Conductivities  and  Viscosities  in  Farmamid. 


19 


all  proportions,  neither  being  soluble  to  any  extent  in  absolute  ether, 
chloroform,  benzol,  hexane,  etc.;  nor  do  they  dissolve  appreciably  the 
aromatic  hydrocarbons,  nitrobenzol,  fats,  oils,  etc.  Further  similarity 
may  be  traced  in  their  solvent  action  on  metallic  salts;  thus,  in  the  cold, 
cobalt  and  nickel  salts  yielding  solutions  in  formamid  similarly  colored 
with  those  in  water,  although  in  some  instances  the  formamid  solutions 
undergo  change  in  color  on  warming,  which  is  probably  due  to  a  predom- 
inance of  the  unionized  salt,  since  formamid  undergoes  a  much  sharper 
decrease  in  association  with  rise  in  temperature  than  water. 

As  Walden  has  pointed  out,  the  similarity  between  water  and  forma- 
mid is  still  closer.  Phosphorus  and  sulphur  are  practically  equally 
insoluble  in  both,  while  iodine  gives  a  brownish-yellow  solution. 
Starch  is  also  soluble  in  formamid,  with  formation  in  concentrated  solu- 
tion of  a  jelly,  and  on  addition  of  formamid  solution  of  iodine  the  starch 
solution  turns  intensely  blue.  The  color,  however,  is  less  permanent 
than  in  water,  owing  to  the  slow  action  of  the  iodine  on  the  solvent. 
The  fluorescent  dyes  also  exhibit  like  phenomena  in  formamid  and  in 
water,  this  being  particularly  marked  in  the  case  of  eosin. 

TABLE  5. 


Physical  constants. 


Water. 


Formamid. 


Molecular  weight 18 

Melting-point 0° 

Boiling-point  760  mm 100° 

Density  0°/4° 9999 

Dielectric  constant 81  (Drude)  "° 

Association  factor  30° ;  3.81  (Ramsay-Shields). 

K-surf ace  tension  20°  to  30° ' 

Minimum  specific  conductivity..!  4X10— 8. 
Average    specific    conductivity      1.5X10—'. 

working  values. 

Viscosity,  25° !  0.00891    (Thorpe    and 

Rodger). 
Dissociation  of  N(CsH6)J  at      i  91  p.  ct. 

V  =  100. 


45 

1.5-2.1  (Walden). 

200-212 

1.151  (Walden,  Davis). 

84  (Walden). 

6.18  (Turner  and  Merry). 

0.65  (Turner  and  Merry). 

2.8X10-'  (Davis  and  Putnam). 

2.7  X10-6  (Davis  and  Putnam). 


0.0324  (Davis). 

93   p.   ct.   (Walden);  98  p. 
(Davis  and  Putnam). 


ct. 


Bruni  and  Mannuelli1  further  show  that  just  as  water  hydrolyzes  the 
salts  of  weak  bases,  such  as  those  of  bismuth  and  antimony,  forming 
unstable  basic  salts;  formamid  by  a  process  of  aminolysis  may  form 
basic  salts  of  these  same  metals;  and  Rohler,2  in  extending  this  work 
has  isolated  characteristic  basic  salts  of  copper,  cobalt,  nickel,  and 
zinc.  He  has  also  obtained  amidates  similar  to  hydrates,  of  which 
PbCl2.  HCONH2  is  an  example,  as  well  as  metal  formamidates  having 
the  general  composition  Me(HNCOH)2.  2HCONH2,  where  Me  may  be 
either  copper,  nickel,  cobalt,  or  zinc. 

In  addition  to  the  above,  Rohler  has  noticed  the  formation  of  well- 
defined  crystalline  compounds  of  formamid  with  the  halogen  acids, 
corresponding  to  the  well-known  mono-,  di-,  and  tri-hydrates. 


'Zeit.  Elektrochem.  11,  554  (1905). 


'Ibid.,  16,  418  (1910). 


20  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

As  is  now  well  established,  the  dielectric  constant  of  the  solvent  is 
a  measure  of  its  own  association,  and  therefore  a  measure  of  its 
dissociating  power  for  electrolytes.  Of  the  common  solvents,  water 
has  the  highest  dielectric  constant,  and  is  the  best  dissociant.  Forma- 
mid,  however,  has  a  higher  dielectric  constant  than  water,  as  will  be 
shown  by  referring  to  the  table  of  physical  constants  of  the  two 
solvents  (table  5),  and  should  therefore  be  a  better  dissociant,  which 
in  our  discussion  will  be  pointed  out  to  be  the  fact. 

PURIFICATION  OF  THE  SOLVENT. 

It  might  be  well  to  state  at  the  outset,  that  neither  a  clean  melting- 
point  nor  a  constant  boiling-point  is  a  sufficient  criterion  for  con- 
ductivity purposes  as  to  the  purity  of  the  solvent,  since  it  has  been 
found  by  ourselves  and  others,  that  a  constant-boiling  liquid  such  as 
formamid,  may  be  separated  into  fractions  of  widely  different  specific 
conductivities.  That  this  is  true  is  due,  no  doubt,  in  this  particular 
case,  to  the  fact  that  minute  quantities  of  the  products  of  hydrolysis, 
such  as  would  have  no  measurable  effect  on  the  apparent  boiling- 
point  of  the  liquid,  on  account  of  the  high  dissociating  power  of  the 
formamid,  produce  a  marked  increase  in  its  conductivity. 

The  chief  criterion,  therefore,  hi  judging  of  the  purity  of  the  solvent 
used  in  this  investigation  was  its  specific  conductivity.  The  material 
with  which  we  started  was  obtained  from  Kahlbaum,  and  had  a  specific 
conductivity  of  about  674  X  10~5,  or  about  that  of  tap-water.  Samples 
obtained  from  Bender  and  Hobein,  from  Schuchardt,  and  from  Hoffman 
and  Kropf  proved  from  the  conductivity  standpoint  to  be  little  or  no 
better  than  the  above. 

As  has  already  been  mentioned,  formamid  is  hygroscopic,  forms  a 
true  solution  with  water,  and  subsequently  undergoes  slow  hydrolysis 
into  ammonium  formate.  The  first  problem,  therefore,  that  must  be 
solved,  was  the  removal  of  any  dissolved  water  not  already  acted  on; 
second,  the  removal  of  the  products  of  hydrolysis  already  present. 
The  method  of  fractional  distillation  made  it  necessary  to  design  and 
construct  suitable  apparatus  for  distillation  in  comparatively  high 
vacuo. 

A  third  problem  presented  itself  in  connection  with  the  preservation 
and  subsequent  manipulation  of  the  solvent  and  of  solutions  in  it,  in 
such  a  manner  as  to  incur  minimum  exposure  to  moisture.  In  addition, 
the  expense  of  the  solvent  made  necessary  the  recovery  of  it  from 
solutions  of  salts  in  this  solvent,  with  the  least  possible  loss  by  decom- 
position. 

The  removal  of  dissolved  water  was  finally  effected  by  the  use  of 
carefully  dehydrated  sodium  sulphate.  After  testing  a  number  of 
other  dehydrating  agents,  such  as  magnesium  sulphate,  calcium  chlor- 
ide, sulphuric  acid  in  vacuo,  etc.,  it  was  found  that  sodium  sulphate 


Conductivities  and  Viscosities  in  Formamid. 


21 


produced  a  smaller  loss  of  materials  from  combination  with  it  than  any 
of  the  other  dehydrating  agents  studied. 

Formamid  was  therefore  allowed  to  stand  for  several  weeks  over 
anhydrous  sodium  sulphate,  in  carefully  sealed,  glass-stoppered  bottles, 
and  placed  in  a  cool,  dark  room.  An  attempt  was  made  to  effect  a 
preliminary  purification  of  formamid  by  fractional  crystallization; 
but  the  end-product,  after  several  fractionations,  invariably  showed  a 
higher  conductivity  than  the  original  substance.  We  were  therefore 
forced  to  conclude,  as  were  Freer  and  Sherman  in  preparing  formamid 
to  study  its  sodium  salts,  that  a  properly  conducted  distillation  was  the 
best  process  available. 


Fio.  6. 

The  apparatus  finally  adopted  for  distillation  is  shown  diagrammati- 
cally  in  figure  5.  By  placing  the  various  parts  in  parallel  rows  it  was 
possible  to  mount  the  whole  apparatus  upon  a  desk-space  of  only  2| 
feet  square,  as  is  shown  in  the  photograph  (plate  1).  This  condensed 
arrangement  of  the  apparatus  given  in  figure  5  is  photographed  in  two 
views  in  plate  1 ,  and  is  the  perfected  apparatus  with  the  exception  of 
the  distillation  head  and  the  position  of  the  manometer. 

In  this  apparatus,  with  all  the  stopcocks  closed,  a  vacuum  of  0.5  mm. 
was  easily  maintained  by  the  Gaede  pump  (fig.  5,  K).  This  pump  was 
mounted,  together  with  a  y ^horsepower  motor  (fig.  5,  L),  controlling 
rheostat  not  shown,  idler  (fig.5,M),  and  switch,  on  a  heavy  maple  wood 
base,  provided  with  carrying  handles  and  rubber  feet. 

During  the  actual  process  of  distillation  the  vacuum  rose  to  from  1.5 
to  2.5  mm.,  since,  on  account  of  the  high  viscosity  of  the  liquid,  it 
was  necessary  to  keep  a  fairly  rapid  current  of  air  flowing  in  through 
the  stopcock  and  drawn-out  portion  of  the  distillation  head  (fig. 5,  E),  in 
order  to  keep  the  liquid  agitated  and  prevent  the  violent  bumping 
which  usually  attended  distillation  in  a  high  vacuum.  The  drying  of 
this  current  of  air  sufficiently,  and  the  proper  method  of  introducing  it, 
proved  to  be  among  the  most  serious  difficulties  encountered.  The 
method  finally  adopted  is  shown  in  figure  5.  The  air,  before  entering 


22  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

the  distillation  flask,  passed  through  a  soda-lime  tower  (not  shown  in 
the  drawing,  but  appears  on  the  left  in  plate  1 ;  then  through  the  two 
wash  bottles  (fig.  5,  A),  and  finally  over  phosphorus  pentoxide  con- 
tained in  the  long  tube  B,  which  was  protected  by  two  glass  stopcocks, 
the  one  sealed  permanently  into  the  tube  and  the  other  securely  fast- 
ened with  sealing-wax  in  the  left-hand  tapered  end  of  B.  This  joint 
made  it  possible  to  open  B  in  renewing  the  phosphorus  pentoxide. 
From  B,  the  current  of  air  passed  into  c,  through  a  short  rubber 
connection  and  a  right-angle  glass  bend,  the  open  end  of  which  is 
connected  with  c  by  a  rubber  joint  made  air-tight  with  rubber  cement. 

The  distillation  head  c  consisted  of  a  short  projecting  tube,  a  well- 
ground  stopcock  of  large  bore;  the  ground  connection  with  the  750 
c.c.  distilling  flask  D,  and  the  long  inner  tube  reaching  to  the  bottom  of 
the  flask  and  drawn  to  a  capillary  at  the  end,  so  as  to  cause  the  air  to 
bubble  in  a  fine  stream  through  the  liquid.  Connection  between  the 
flask  and  the  condenser  was  effected  by  a  ground  joint  designed  to 
prevent  the  liquid  distilling  over  from  coming  in  contact  with  the  stop- 
cock grease  used  in  making  the  joint  air-tight.  The  condenser  was  also 
ground  into  the  Axtell  receiver  F,  which  is  too  well  known  to  require 
further  description,  other  than  to  state  that  the  capacity  of  the  barrel 
was  about  100  c.c.  and  that  of  the  small  bulb  250  c.c.  The  receiver 
allowed  the  removal  of  different  fractions  without  interrupting  the 
vacuum. 

The  receiver  was  connected  with  the  pump  K,  through  the  gas  wash 
bottles  A',  the  soda-lime  tower  i,  and  the  three-way  stopcock  J.  One 
of  the  wash  bottles  was  filled  just  to  the  upper  level  of  the  holes  in 
the  inner  cylinder  with  sulphuric  acid;  the  other  was  empty,  serving 
as  a  trap  in  the  case  of  back-pressure.  The  construction  of  the  wash 
bottles  secured  a  maximum  exposure  of  acid  to  the  possible  alkaline 
vapors  (NH3)  without  appreciable  back-pressure. 

Liquid  formamid  was  introduced  into  the  flask  at  any  time  during  the 
distillation  without  destroying  the  vacuum,  by  closing  the  stopcock  on 
c,  temporarily  disconnecting  the  air  inlet,  and  attaching  at  the  same 
point  the  flask  and  siphon  N.  On  opening  the  stopcock  the  diminished 
pressure  in  the  flask  drew  in  any  liquid  in  N. 

During  the  distillation  the  flask  was  immersed  in  a  bath  of  hardened 
cottonseed  oil  ("Crisco"),  into  which  the  thermometer  dipped.  Ordinary 
rectified  cottonseed  oil  was  first  used,  but  this  was  soon  discontinued, 
since  the  oil  quickly  became  rancid  and  offensive  when  heated.  The 
"Crisco,"  however,  remains  sweet  and  untainted  even  after  continued 
use.  The  temperature  in  the  outer  bath  was  maintained  at  from 
10°  to  15°  higher  than  the  boiling-point  of  the  liquid  at  a  given  pressure, 
the  usual  temperature  being  100°  to  115°.  In  the  apparatus  used  in  the 
preliminary  work  a  thermometer  was  sealed  into  the  air-inlet  tube  of 
c,  with  its  bulb  well  down  in  the  flask,  and  the  boiling-point  of  the 


I.     Distilling  Appari 


2.     Spectroscope  with  Cover. 


Conductivities  and  Viscosities  in  Formamid. 


23 


liquid  was  observed;  but  this  thermometer  was  subsequently  discarded 
and  the  temperature  controlled  solely  by  the  thermometer  in  the  outer 
bath,  so  as  to  obtain  about  100  c.c.  of  distillate  every  15  minutes. 

As  a  rule,  starting  with  Kahlbaum's  so-called  c.  p.  formamid,  from 
three  to  four  distillations  yielded  a  product  sufficiently  pure  for  our 
purpose.  The  decrease  in  the  conductivity  with  successive  distillation 
is  shown  in  table  6: 

TABLE  6. — Specific  Conductivity. 
[Original  material,  674.0X10-'.] 


First 
distillation. 

Second 
distillation. 

Third 
distillation. 

First  fraction  

785.8  X  10  -5 

204.2  X10-6 

24  6X10-S 

Second  fraction  

435.0  X10-6 

153.8  X10-5 

10.3X10-' 

Third  fraction  

112.3X10-' 

67.4  X  10-' 

4.8X10-' 

While  the  end-fraction  in  this  case  was  sufficiently  pure  for  use, 
a  fourth  distillation  yielded  a  product  with  a  conductivity  of  about 
3-4  X  10~6.  Although  the  exact  minimum  conductivity  of  this  solvent 
has  thus  far  not  been  ascertained,  we  have  obtained  small  lots  with  a 
specific  conductivity  of  2.8  XlO"6  to  4.7X10"6,  or  from  2  to  3  times 
the  value  for  the  conductivity  water  used  in  this  laboratory.  Walden 
used  in  his  work  a  product  with  a  conductivity  of  7.5X10"4,  but 
obtained  a  small  fraction  of  about  10  per  cent  of  the  original  material, 
with  a  conductivity  of  4.7X10"8.  Our  solvent  was  from  35  to  40 
per  cent  of  the  original  volume  with  a  maximum  specific  conductiv- 
ity of  4.8X10"6,  with  a  minimum  of  1.37X10"6  and  an  average  of 
2.7X10-5. 

SALTS. 

On  account  of  the  instability  of  the  solvent  formamid  in  the  presence 
of  moisture,  the  salts  used  in  this  investigation  were  all  dehydrated  with 
special  care  at  the  highest  temperature  which  it  was  possible  to  use  for 
these  substances. 

SOLUTIONS. 

The  more  concentrated  solutions  in  formamid  were  made  up  by  direct 
weighing;  the  more  dilute  by  diluting  the  more  concentrated.  This 
operation  was  much  facilitated  by  the  use  of  two  burettes  holding  50 
and  10  cm.  respectively;  the  one  being  employed  for  the  solvent,  the 
other  for  the  one-tenth  normal  solution.  Each  burette  was  connected 
with  the  reservoir  containing  the  solvent  or  the  solution,  by  a  siphon 
provided  with  a  glass  stopcock,  access  of  moisture  being  prevented 
by  calcium  chloride  tubes  connected  with  both  the  burettes  and  the 
reservoirs. 


24 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


The  solvent  was  kept  in  half-liter  glass-stoppered  Erlenmeyer  flasks, 
and  the  solutions  when  made  up  were  preserved  in  35  c.c.  flasks  of 
similar  design,  which  were  sealed  with  rubber  cement. 

On  account  of  the  high  price  of  the  solvent,  only  25  c.c.  of  each 
solution  was  prepared,  this  amount  serving  both  for  conductivity  and 
viscosity  measurements.  All  operations  in  preparing  the  solutions  were 
carried  out  at  20°. 

APPARATUS. 

Conductivity  apparatus. — The  conductivity  apparatus  was  identical 
with  that  employed  in  our  earlier  work  on  binary  and  ternary  mixtures 
of  water,  acetone,  and  glycerol;  the  method  of  obtaining  duplicate 
readings  and  other  details  being  exactly  the  same  as  in  the  earlier  work, 
excepting  the  use  of  a  rocking  commutator  with  mercury  contacts 
instead  of  the  two-blade,  double-throw  switch  used  in  reading  both 
ends  of  the  bridge,  the  object  being  to  eliminate  as  nearly  as  possible 
all  external  resistance.  The  conductivity  cells  were  of  the  type  recently 
employed  in  this  laboratory  for  work  with  non-aqueous  solvents,  and 
were  carefully  calibrated  at  regular  intervals. 

Viscosity  apparatus. — This  was  essentially  the  same  as  in  our  earlier 
investigations.  We  have  designed  and  used  an  improved  support  for 
the  viscosimeters,  which  is  particularly  well  adapted  to  our  new  thermo- 
stat. A  new  pyknometer  has  also  been  devised,  which  has  proved  to 
have  advantages  over  the  older  form.  (See  Chapter  I.) 

Thermostats. — The  new  form  of  constant-temperature  bath  for  con- 
ductivity and  viscosity  investigations  has  already  been  fully  described 
in  Chapter  VI  of  Publication  No.  210  of  the  Carnegie  Institution  of 
Washington. 

DISCUSSION  OF  RESULTS. 

Tables  8  to  27,  inclusive,  give  the  molecular  conductivity,  dissocia- 
tion, viscosity,  and  fluidity,  as  well  as  the  temperature  coefficients 
both  of  conductivity  and  fluidity,  for  all  of  the  salts  studied.  Measure- 
ments were  made  at  15°,  25°,  and  35°. 

TABLE  7. — Comparison  of  the  Various  Solvents. 


Solvent. 

Dielectric 
constant. 

Association 
factor. 

Viscosity 
at  25°. 

HCN 

95 

HCONHj 

94 

6  18 

0  0326 

H2O  

81  7 

3  81 

0  0089 

CHjOH 

32  5 

3  43 

0  0056 

CjHsOH  .  . 

21  7 

2  74 

0  0111 

C»H6  (OH)i  

16  5 

1  80 

5  854 

CH.-CO-CH,  

20.7 

1.26 

0.0035 

Conductivities  and  Viscosities  in  Formamid. 


25 


TABIE  8. — Sodium  Bromide  in  Formamid. 
[K.'for  15°  =1.53X10-';  for  25°  =1.99X10-';  for  35°  =2.46X10-']. 


Temperature  coefficients. 

Molecular,  conductivity 

Dissociation. 

Per  cent. 

Conductivity 

units. 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15  to  25° 

25  to  35° 

2 

11.83 

15 

.53 

19.76 

62.1 

63.1 

63.4 

0.0313 

0.0272 

0.370 

0.423 

4 

13.80 

18 

.01 

22.82 

72.4 

73.2 

73.6 

0.0305 

0.0267 

0.421 

0.481 

10 

16.20 

21 

.09 

26.64 

85.0 

85.7 

86.0 

0.0302 

0.0263 

0.489 

0.555 

50 

17.46 

22 

.62 

28.55 

91.6 

91.9 

92.1 

0.0296 

0.0262 

0.516 

0.593 

200 

18.79 

24 

.36 

30.59 

98.6 

98.9 

98.7 

0.0296 

0.0256 

0.557 

0.623 

400 

19.06 

24 

.62 

30.99 

100.0 

100.0 

100.0 

0.0292 

0.0259 

0.0556 

0.637 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

ij!5° 

1725° 

ij  35°           *>  15°         f  25° 

P35° 

15  to  25° 

25  to  35° 

2 

0.05578 

0.04081 

0.03120        17.93       24.50 

32.05 

0.0367 

0.0308 

4 

0.04883 

0.03678 

0.02837       20.48       27.19 

35.25 

0.0328 

0.0296 

10 

0.04523 

0.03413 

0.02666       22.11        29.30 

37.51 

0.0325 

0.0280 

Solv. 

0.04274 

0.03194 

0.02511        23.40       31.31 

39.83 

0.0338 

0.0272 

*K  is  the  specific  conductivity  of  the  solvent. 

TABLE  9. — Sodium  Iodide  in  Formamid. 
[K,  for  15°  =2. 23X10-';  for  25°  =2. 83X10-';  for  35°  =3.48X10"'.] 


Molecular  conductivity 

Dissociation. 

Per  cent. 

Conductivity 
units. 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15  to  25° 

25to35° 

2 

11.92 

15.71 

20.01 

64.7 

65.6 

66.1 

0.0318 

0.0274 

0.379 

0.430 

4 

14.17 

18.57 

23.59 

76.9 

77.5 

78.0 

0.0311 

0.0270 

0.440 

0.502 

10 

15.91 

20.67 

26.05 

86.3 

86.3 

86.1 

0.0299 

0.0260 

0.476 

0.538 

50 

17.17 

22.33 

28.17 

93.2 

93.2 

93.1 

0.0301 

0.0262 

0.516 

0.584 

200 

18.27 

23.80 

29.98 

99.1 

99.3 

99.1 

0.0303 

0.0260 

0.553 

0.618 

400 

18.43 

23.96 

30.25 

100.0 

100.0 

100.0 

0.0300 

0.0263 

0.553 

0.629 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

7/15° 

j)25° 

17  35°           v  15° 

P25° 

P35° 

15  to  25° 

25  to  35° 

2 

0.05425 

0.03997 

0.03065       18.43 

25.02 

32.63 

0.0357 

0.0304 

4 

0.04822 

0.03640 

0.02800       20.74 

27.47 

35.71 

0.0325 

0.0300 

10 

0.04532 

0.03381 

0.02653       22.07 

29.58 

37.69 

0.0340 

0.0274 

Solv. 

0.04307 

0.03302 

0.02570       23.22 

30.29 

38.91 

0.0304 

0.0285 

26  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

TABLE  10. — Sodium  Chromale  in  Formamid. 
[K,  for  15°=1.30X10-5;  for  25°  =1.65X10-";  for  35°  =2.0.3  XIO"5.] 


V 

Molecular  conductivity 

Dissociation. 

Per  cent. 

Conductivity 

units. 

15° 

25° 

. 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

1 
15  to  25°  25  to  35° 

10 
60 
200 
400 
800 
1,600 

24.57 
31.46 
34.34 
39.00 
42.64 
72.06 

33.66 
41.48 
44.82 
50.79 
55.07 
93.51 

42.66 
52.18 
56.38 

0.0316 
0.0319 
0.0305 
0  0302 

0.0267 
0.0258 
0.0258 

0.809     0.900 
1.002      1.070 
1.048      1.156 
1   179  1 

69.16 
116.44 

0.0292 
0.0298 

0.0256 
0.0245 

1  .  243      1   409 
2.145     2.293 

V 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

.,15° 

>)250 

7j35° 

P15° 

«.25° 

^35° 

15  to  25° 

25  to  35° 

10 
Solv. 

0.04966 
0.04301 

0.03633 
0.03221 

20.14 
23.25 

27.30 
31.05 

TABLE  11. — Potassium  Chloride  in  Formamid. 
[K,  for  15°=2.94X10-6;  for  25°  =3.75X10-*;  for  35°  =4.69  XIO"'.] 


Molecular  conductivity 

Dissociation. 

Per  cent. 

Conductivity 
units. 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15  to  25° 

25  to  35' 

2 

14.12 

18 

25 

22.94 

67.9 

68.3 

68.7 

0.0292 

0.0257 

0.413 

0.469 

4 

16.15 

20.84 

26.13 

77.6 

78.0 

78.3 

0.0290 

0.0252 

0.469 

0.529 

10 

18.06 

23 

27 

28.99 

86.8 

87.1 

86.8 

0.0288 

0.0246 

0.521 

0.572 

50 

19.27 

24 

94 

31.12 

92.6 

93.3 

93.2 

0.0294 

0.0248  |  0.567 

0.618 

200 

20.66 

26 

60 

33.10 

99.3 

99.5 

99.1 

0.0288 

0.0244      0.594 

0.650 

400 

20.80 

26 

73 

33.39 

100.0 

100.0 

100.0 

0.0285 

0.0249 

0.593 

0.666 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

r,  15° 

7j25° 

1 
7)  35°           v  15°    I     v  25°         v  35° 

15  to  25° 

25  to  35° 

I                 i 

2 

0.05251 

0.03922 

0.03000        19.04       25.50  j     33.33 

0.0339 

0.0307 

4 

0.04724 

0.03572 

0.02794       21.17       28.00       35.79 

0.0323 

0.0278 

10 

0.04457 

0.03386 

0.02642        21.94       29.53  !     37.85 

0.0338 

0.0282 

Solv. 

0.04304 

0.03256 

0.02542       23.23       30.71  j     39.34 

! 

0.0322 

0.0281 

Conductivities  and  Viscosities  in  Formamid. 


27 


TABLE  12. — Potassium  Iodide  in  Formamid. 
[K,  for  15°  =2. 23X10-';  for  25°=2. 83X10-*;  for  35°  =3.48X10~'.] 


Molecular  conductivity 

Dissociation. 

Per  cent. 

Conductivity 
units. 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15to25° 

25  to  35° 

2 

14.42 

18 

67 

23.29 

70.7 

71.0 

71.0 

0.0295 

0.0247 

0.425 

0.462 

4 

16.32 

21 

05 

26.34 

80.0 

80.1 

80.3 

0.0290 

0.0251 

0.473 

0.529 

10 

18.23 

23 

45 

29.21 

89.4 

89.2 

89.0 

0.0286 

0.0246 

0.522 

0.576 

50 

19.25 

24 

87 

31.03 

94.4 

94.6 

94.5     0.0292 

0.0248 

0.562 

0.616 

200 

20.39 

26 

28 

32.81 

100.0 

100.0 

100.0     0.0289 

0.0248 

0.589 

0.653 

400 

20.26 

?6 

m 

32.82 

1 

!/• 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

ij  15° 

•n  25° 

Jt  35°           v  15°         f  25° 

»>35° 

15  to  25° 

25  to  35° 

2 

0.04982 

0.03710 

0.02884       20.07    

34  67 

0  0343 

0  0286 

4 

0.04631 

0.03525 

0.02716       21.59       28.37 

36.81 

0.0314 

0.0298 

10 

0.04418 

0.03353 

0.02629       22.64       29.82 

38.04 

0.0318 

0.0274 

Solv. 

0.04307 

0.03302 

0.02570        23.22        30.29 

38.91 

0.0304 

0.072 

TABLE  13. — Potassium  Sulphocyanate  in  Formamid. 
[K,  for  15°  =1.07X10-*;  for  25°  =1.37  X1Q-6;  for  35°  =1.69X10-'.] 


Molecular  conductivity 

Dissociation. 

Per  cent. 

Conductivity 

units. 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15  to  25° 

25  to  35° 

2 

15.15 

19 

.10 

24.28 

70.0 

69.2 

70.4 

0.0261 

0.0271 

0.395 

0.518 

4 

17.42 

22 

.31 

27.87 

80.6 

80.8 

80.8 

0.0281 

0.0249 

0.489 

0.556 

10 

19.04 

24 

38 

30.35 

88.0 

88.3 

88.0 

0.0280 

0.0245 

0.534 

0.597 

50 

20.02 

25 

61 

32.16 

92.6 

92.8 

93.3 

0.0279 

0.0256 

0.558 

0.655 

200 

21.21 

27.19 

33.89 

98.1 

98.5 

98.3 

0.0282 

0.0246 

0.598 

0.670 

400 

21.62 

27 

60 

34.47 

100.0 

100.0 

100.0 

0.0277 

0.0249 

0.598 

0.687 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

1715° 

7,25° 

i)35° 

v  15°         v  25° 

V35° 

15  to  25° 

25  to  35° 

2 

0.04891 

0.03657 

0.02838 

20.45       27.35 

35.24 

0.0337 

0.0287 

4          0.045S5 

0.03473 

0.02713 

21.81        28.79 

36.86 

0.0320 

0.0280 

10          0.04369 

0.03280 

0.02574 

22.89       30.49 

38.85 

0.0332 

0.0274 

Solv.     \     0  .  04294 

0.03258 

0.02554 

23.29       30.69 

39.15 

0.0314 

0.0276 

28 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


TABLE  14. — Ammonium  Bromide  in  Formamid. 
[K,  for  15°  =2. 94X10-';  for  25°  =3.76X10~5;  for  35°  =4.67X10-'.] 


Molecular  conductivity 

Dissociation. 

Per  cent. 

Conductivity 
units. 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15  to  25° 

25  to  35° 

2 

16.31 

20 

.85 

25.79 

70.6 

70.6 

70.0 

0.0278 

0.0237 

0.454 

0.494 

4 

18.21 

23 

.31 

28.88 

78.8 

78.9 

78.3 

0.0280 

0.0247 

0.510 

0.557 

10 

20.57 

26 

.21 

32.53 

89.0 

88.7 

88.3 

0.0274 

0.0241 

0.564 

0.632 

50 

21.82 

27.94 

34.82 

94.4 

94.6 

94.5 

0.0280 

0.0246 

0.612 

0.688 

200 

22.63 

28 

.83 

35.93 

97.9 

97.6 

97.5 

0.0274 

0.0246 

0.620 

0.710 

400 

23.11 

29.54 

36.86 

100.0 

100.0 

100.0 

0.0278 

0.0248 

0.643 

0.732 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

i/15° 

1725° 

7)  35°           v  15° 

P25° 

«,35° 

15  to  25° 

25  to  35° 

2 

0.04795 

0.03607 

0.02776       20.86 

27.73 

36.02 

0.0330 

0.0299 

4 

0.04550 

0.03455 

0.02680       21.98 

28.94 

37.31 

0.0317 

0.0289 

10 

0.04399 

0.03273 

0  .  02635       22  .  73 

30.55 

37.95 

Solv. 

0.04304 

0.03256 

0.02542       23.23 

30.71 

39.34 

0.0322 

0.0281 

TABLE  15. — Ammonium  Iodide  in  Formamid. 
[K,  for  15°  =3. 68X10-';  for  25°  =4.71  X1Q-6;  for  35°  =5.87X10~6.] 


Molecular  conductivity 

Dissociation. 

Per  cent. 

Conductvity 
units. 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15to25° 

25  to  35° 

2 

16.90 

21 

.70 

26.80 

73.6 

73.1 

72.5 

0.0284 

0.0281 

0.480 

0.610 

4 

18.78 

24 

.06 

29.88 

81.8 

81.0 

80.8 

0.0281 

0.0240 

0.528 

0.582 

10 

20.48 

26 

.32 

32.44 

89.2 

88.6 

87.7 

0.0285 

0.0233 

0.584 

0.612 

60 

21.31 

27.39 

33.92 

92.8 

92.3 

91.2 

0.0285 

0.0238 

0.608 

0.653 

100 

22.20 

28 

.40 

35.40 

96.7 

95.7 

95.8 

0.0279 

0.0246 

0.620 

0.700 

400 

22.96 

29 

.69 

36.97 

100.0 

100.0 

100.0 

0.0293 

0.0245 

0.673 

0.728 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

i;150 

ij25° 

1735° 

v  15°        v  25° 

«>35° 

15  to  25° 

25  to  35° 

2 

0.05091 

0.03856 

0.03007 

19.64       25.93 

33.26 

0.0320 

0.0282 

4 

0.04571 

0.03417 

0.02669 

21.88       29.27 

37.47 

0.0292 

0.0280 

10 

0.04367 

0.03311 

0.02607 

22.90       30.20 

38.36 

0.0319 

0.0270 

Solv. 

0.04201 

0.03207 

0.02496 

23.80       31.18 

40.06 

0.0310 

0.0285 

Conductivities  and  Viscosities  in  Formamid. 


29 


TABLE  l&.—Telramethylammonium  Chloride  in  Formamid. 
,K,  for  15°=1.53X10-6;  for  25°  =1 .99X10-';  for  35°  =2.46 X1Q-5.] 


Molecular  conductivity 

Dissociation. 

Per  cent. 

Conductivity 
units. 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15to25° 

25  to  35° 

2 

14.62 

18.72 

23.38 

68.6 

68.5 

68.1 

0.0280 

0.0249 

0.410 

0.466 

4 

16.92 

21 

80 

27.27       79.4 

79.7 

79.4 

0.0288 

0.0251 

0.488 

0.547 

10 

18.37 

23 

65 

29.32  j     86.2 

86.5 

85.4 

0.0287 

0.0240 

0.528 

0.567 

50 

19.64 

25 

32 

31.65       92.1 

92.6 

92.1 

0.0289 

0.0250 

0.568 

0.633 

200 

21.19 

27 

34 

34.11 

99.5 

100.0 

99.3 

0.0290 

0.0248 

0.615 

0.677 

400 

21.30 

27 

34 

34.35 

100.0 

100.0 

100.0 

0.0283 

0.0256 

0.604 

0.701 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

7/25° 

i?  35°           f  15°         f  25" 

,350 

15  to  25° 

25  to  35° 

2 

0.04775 

0.03578 

0.02791       20.94       27.95 

35.83 

0.0335 

0.0282 

4 

0.04503 

0.03427 

0.02679       22.21        29.18 

37.33 

0.0314 

0.0279 

10         0.04394 

0.03313 

0.02601        22.76       30.18 

38.45 

0.0326 

0.0274 

Solv.         0.04274 

0.03194 

0.02511        23.40       31 

.31 

39.83 

0.0338 

0.0272 

TABLE  17. — Tetraethylammonium  Iodide  in  Formamid. 
[K,  for  15°=2.47X10-S;  for  25°=3. 17X10-6;  for  35°  =3.94X10-'.] 


Molecular  conductivity. 

Dissociation. 

Per  cent. 

Conductivity 
units. 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15  to  25° 

25  to  35° 

2 

11.78 

15.42 

19.47 

62.6 

63.5 

64.0 

0.0309 

0.0262 

0.364 

0.405 

4 

13.86 

18 

10 

22.78 

73.7 

74.5 

74.9 

0.0306 

0.0259 

0.424 

0.468 

10 

15.74 

20 

31 

25.31 

83.7 

83.6 

83.2 

0.0290 

0.0246 

0.457 

0.500 

50 

17.60 

22 

69 

28.47 

93.6 

93.5 

93.6 

0.0298 

0.0255 

0.509 

0.578 

100 

18.45 

23 

79 

29.85 

98.1 

98.0 

98.1 

0.0289 

0.0255 

0.534 

0.606 

200 

18.81 

24 

28 

30.43 

100.0 

100.0 

100.0 

0.0291 

0.0253 

0.547 

0.615 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

i)15° 

i)25° 

1735° 

«,15° 

«,25° 

*>35° 

15  to  25° 

25  to  35° 

2 

0.04879 

0.03618 

0.02802 

20.50 

27.64 

35.68 

0.0349 

0.0291 

4 

0.04573 

0.03486 

0.02687 

21.87 

28.69 

37.22 

0.0312 

0.0297 

10 

0.04431 

0.03336 

0.02607 

22.57 

29.98 

38.36 

0.0328 

0.0280 

Solv. 

0.04284 

0.03256 

0.02561 

23.34 

30.71 

39.05 

0.0316 

0.0271 

30 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


TABLE  18.—  Rubidium  Chloride  in  Formamid. 
[K,  for  15°  =2. 47X10-';  for  25°  =3. 17X1Q-';  for  35°  =3.94  X10~'.] 


Molecular  conductivity.            Dissociation. 

Per  cent. 

Conductivity 
units. 

V 

15° 

25° 

35°     i     15° 

25° 

35° 

15  to  25° 

25  to  35° 

15  to  25° 

25  to  35° 

2 

14.28 

18.46 

23.10       67.7 

68.1 

68.6 

0.0292 

0.0251 

0.418 

0.464 

4 

16.31 

21 

00 

26.17       77.3 

77.5 

78.0 

0.0288 

0.0251 

0.469 

0.527 

10 

18.77 

24 

09 

30.10 

89.0 

88.9 

89.3 

0.0283 

0.0249 

0.532 

0.601 

50 

19.84 

25 

54 

31.82 

94.1 

94.3 

94.4 

0.0287 

0.0246 

0.570 

0.628 

100 

20.59 

26 

49 

33.03 

97.6 

97.8 

98.0 

0.0287 

0.0247 

0.590 

0.654 

200 

21.09 

27.08 

33.69 

100.0 

100.0 

100.0 

0.0284 

0.0244 

0.599 

0.661 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

i)  15° 

7(25° 

1;  35°           v  15°         v  25° 

V35° 

15  to  25° 

25  to  35° 

2 

0.05246 

0.03869 

0.02961        19.06       25.85 

33.77 

0.0356 

0.0308 

4 

0.04773 

0.03606 

0.02778       20.95       27.73 

36.00 

0.0324 

0.0298 

10 

0.04486 

0.03396 

0.02623       22.29       29.45 

38.12 

0.0321 

0.0295 

Solv. 

0.04284 

0.03256 

0.02561        23.34       30.71 

39.05 

0.0316 

0.0271 

TABLE  19. — Rubidium  Bromide  in  Formamid. 
[K,  for  15°=4.79X10~6;  for  25°  =6. 07X10"';  for  35°  =7.62X10~S.] 


Molecular  conductivity 

Dissociation. 

Per  cent. 

Conductivity 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15  to  25° 

25  to  35° 

4 

16.36 

21 

26 

26.52 

79.8 

81.1 

81.2 

0.0300 

0.0247 

0.490 

0.526 

10 

18.47 

23 

79 

29.63 

90.1 

90.8 

90.7 

0.0286 

0.0245 

0.532 

0.584 

50 

20.49 

26 

21 

32.68 

100.0 

100.0 

100.0 

0.0279 

0.0247 

0.572 

0.647 

100 

20.36 

?6 

87 

32  71 

1 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

ijl5° 

ij25° 

1»35° 

«>15° 

P25°' 

«>35° 

15  to  25° 

25  to  35° 

2 

Solution 

supersaturat 

ed  at  25°. 

4 

0.04661 

0.03501        0.02717 

21.46 

28.56 

36.81        0.0331 

0.0289 

10 

0.04462 

0.03394 

0.02648 

22.41 

29.46 

37.76       0.0315 

0.0282 

Solv. 

0.04312 

0.03260 

0.02564 

23.19 

30.68 

39.00 

0.0323 

0.0271 

Conductivities  and  Viscosities  in  Formamid. 

TABLE  20. — Rubidium  Iodide  in  Formamid. 
[K,  for  15°=3.68X10-S;  for  25°  =4.71  X10~5;  for  35°  =5.87 XIO"5.] 


31 


Molecular  conductivity 

Dissociation. 

Per  cent. 

Conductivity 
units. 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15  to  25° 

25  to  35° 

2 

14.73 

19 

01 

23.78 

72.4 

72.0 

72.2 

0.0290 

0.0251 

0.428 

0.477 

4 

16.87 

21 

71 

27.21 

82.9 

82.2 

82.6 

0.0287 

0.0253 

0.484 

0.550 

10 

18.60 

24 

00 

29.83 

91.4 

90.9 

90.5 

0.0290 

0.0243 

0.540 

0.583 

50 

19.59 

25 

27 

31.49 

96.3 

95.7 

95.6 

0.0290 

0.0246 

0.568 

0.622 

100 

20.23 

26 

06 

32.52 

99.5 

98.7 

98.7 

0.0288 

0.0248 

0.568 

0.646 

200 

20.34 

26 

41 

32.95 

100.0 

100.0 

100.0 

0.0298 

0.0248 

0.607 

0.654 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

7i  15° 

7[  25° 

i)35° 

«,15° 

«,25° 

«.35° 

15  to  25° 

25  to  35° 

2 

0.04945 

0.03516 

0.02865 

20.22 

27.66 

34.90 

0.0368 

0.0265 

4 

0.04643 

0.03503 

0.02737 

21.54 

28.55 

36.54 

0.0325 

0.0280 

10 

0.04432 

0.03348 

0.02634 

22.56 

29.87 

37.97 

0.0324 

0.0271 

Solv. 

0.04201 

0.03207 

0.02496 

23.80 

31.18 

40.06  !     0.0310 

0.0285 

TABLE  21. — Rubidium  Nitrate  in  Formamid. 
[K,  for  15°=2.47X10-5;  for  25°  =3.17 X1Q-5;  for  35°  =3.94X10-'.] 


Molecular  conductivity. 

Dissociation. 

Per  cent. 

Conductivity 
units. 

1 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25°  25  to  35°  15  to  25° 

25  to  35° 

4 

16.20 

20 

79 

25.85 

77.7 

77.6 

77.6 

0.0283      0.0243 

0.459 

0.506 

10 

18.85 

24 

07 

30.00 

90.4 

89.8 

90.1 

0.0277  I  0.0246 

0.522 

0.593 

50 

20.22 

25 

97 

32.23 

97.0 

96.9 

96.8 

0.0284      0.0241 

0.575 

0.626 

100 

20.85 

26 

80 

33.31 

100.0 

100.0 

100.0 

0.0285     0.0243 

0.595 

0.651 

200 

90  76 

°6 

68 

33  02 

!             '             1 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

V 

,15° 

7)25° 

7,  35°            v  15°         v  25° 

,35° 

15  to  25° 

25  to  35° 

4 

0.04586 

0 

03430 

0  02673        21.81        29.16 

37.41 

0 

0337 

0.0283 

10 

0.04417 

0 

03346 

0.02622        22.64        29.89 

38.15 

0 

0320 

0.0276 

Solv. 

0.04284 

0 

03256 

0.02561        23.34        30.71 

39.05 

0 

0316 

0.0271 

32  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

TABLE  22. — Caesium  Chloride  in  Formamid. 
[K,  for  15°  =  1.34X10-5;  for  25°  =1.72 XIO"6;  for  35°  =2. 10X10-6.] 


V 

Molecular  conductivity. 

Dissociation. 

Per  cent. 

hi 
n 

ctivity 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15  to  25° 

25  to  35° 

4 
10 
200 
400 

16.83 
18.75 
21.40 
21.06 

21.69 
24.09 
27.54 
27.15 

27.15 
30.03 
34.49 
34.11 

78.6 
87.6 
100.0 

78.8 
87.5 
100.0 

78.7 
87.1 
100.0 

0.0289 
0.0285 
0.0287 
0.0289 

0.0252 
0.0247 
0.0252 
0.0256 

0.486 
0.534 
0.614 
0.609 

0.546 
0.594 
0.695 
0.696 

V 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

r/150 

i;250 

1735° 

<P  15°         f  25° 

,,35° 

15  to  25° 

25  to  35° 

4 
10 
Solv. 

0.04735 
0.04481 
0.04317 

0.03578 
0.03395 
0.03245 

0.02789 
0.02654 

21. 
22. 
23. 

12       27.95 
32       29.46 
16       30.81 

35.86 
37.68 

0.0323 
0.0320 
0.0000 

0.0283 
0.0279 

TABLE  23.  —  Caesium  Nitrate  in  Formamid. 
[K,  for  15°=1.34X10-6;  for  25°  =1.72  X1Q-6;  for  35°  =2.10X10-'.] 

V 

Molecular  conductivity 

Dissociation. 

Per  cent. 

Conductivity 
units. 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15to25° 

25  to  35° 

4 
10 

100 
200 
400 

17.00 
18.90 
21.36 
21.75 
21.59 

21.79 
24.23 
27.42 
27.90 
27.64 

27.10 
29.92 
33.99 
34.61 
34.43 

78.2 
86.9 
98.2 
100.0 

78.1 
86.8 
98.3 
100.0 

78.3 
86.4 
98.2 
100.0 

0.0282 
0.0282 
0.0284 
0.0283 
0.0280 

0.0244 
0.0235 
0.0240 
0.0241 
0.0246 

0.479 
0.533 
0.606 
0.615 
0.605 

0.531 
0.569 
0.657 
0.671 
0.679 

V 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

7jl5° 

))25° 

7735° 

ip  15°         <p  25° 

,35° 

15  to  25° 

25  to  35° 

4 
10 

Solv. 

0.04614 
0.04456 
0.04317 

0.03478 
0.03362 
0.03245 

0.02725 
0.02632 

21. 
22. 
23. 

67       28.75 
44        29.74 
16       30.81 

36.70 
37.9 

0.0327 
0.0325 

0.0276 
0.0277 

Conductivities  and  Viscosities  in  Formamid. 
;  24. — Lithium  Nitrate  in  Formamid. 


33 


V 

Molecular  conductivity. 

Dissociation. 

Per  cent. 

Conductivity 
units. 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15  to  25° 

25  to  35° 

2 
4 
10 
50 
200 
400 

12.16 
14.10 
16.11 
17.38 
18.42 
18.23 

15 
18 
20 
22 
23 
23 

57 

07 
58 
29 
66 
63 

66.0 
76.5 
87.5 
94.4 
100.0 

65.8 
76.4 
87.0 
94.2 
100.0 

86.9 
94.1 
99.8 
100.0 

0.0280 
0.0281 
0.0277 
0.0283 
0.0284 
0.0296 

0.0241 
0.0242 
0.0240 
0.0244 

0.341 
0.397 
0.447 
0.491 
0.524 
0.540 

0.496 
0.539 
0.568 
0.577 

25.54 
27.68 
29.34 
29.40 

V 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

i)15° 

ij25° 

17  35°           v  15°        v  25° 

P35° 

15  to  25° 

25  to  35° 

2 
4 
10 
Solv. 

0.05191 
0.04720 
0.04460 
0.04272 

0.03873 
0.03571 
0.03157 
0.03196 

0.03019        19.26       25.82 
0.02786       21.19       28.00 
0.02646       22.42       31.68 
0.02503       23.41        31.29 

33.12 
35.89 
37.79 
39.95 

0.0340 
0.0321 

0.0283 
0.0282 

0.0337 

0.0277 

TABLE  25. — Barium  Chloride  in  Formamid. 
[K,  for  15°=1.58X10-5;  for  25°=2.02X10-8;  for  35°  =2.45 XIQ-6.] 


Molecular  conductivity 

Dissociation. 

Per  cent. 

Conductivity 
units. 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15to25° 

25  to  35° 

10 

30.58 

40.00 

50.12 

70.3 

71.6 

70.1 

0.0308 

0.0253 

0.942 

1.012 

50 

38.06 

49 

43  i  61.90 

87.5 

88.5 

86.6 

0.0299 

0.0234 

1.137 

1.157 

200 

40.38 

52 

48  |  65.93 

92.9 

93.9 

92.2 

0.0300 

0.0256 

1.210 

1.345 

800 

43.45 

55 

82 

70.38 

99.9 

99.9 

98.5 

0.0285 

0.0261 

1.237 

1.456 

1,600 

43.48 

55 

86 

71.48 

100.0 

100.0 

100.0 

0.0285 

0.0280 

1.238 

1.562 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

«jl5° 

i;25° 

7)35° 

v  15°         v  25° 

P35° 

15  to  25° 

25  to  35" 

10 

0.04941 

0.03702 

0.02878 

20.24       27.01 

34.75 

0.0335 

0.028 

Solv. 

0.04301 

0.03221 

0.02510 

23.25       31.05 

39.84 

0.0335 

0.0283 

34  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

TABLE  26. — Mercuric  Chloride  in  Formamid. 
[K,  for  15°=1.30X10-5;  for  25°  =  1.65X10-6;  for  35°  =2.03 Xlf)-'.] 


V 

Molecular  conductivity. 

Dissociation. 

Per  cent. 

Conductivity 
units. 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15  to  25° 

25  to  35° 

4 
10 
50 
200 
400 
1,000 
1,600 
3,200 

0.43 
0.63 
1.09 
1.74 
2.73 

0 
0 
1 
2 
4 

n<i 

59 
87 
56 
53 
15 
•19 

0.83 
1.27 
2.22 
3.68 
5.82 

0.0372 
0.0381 
0.0431 
0.0437 
0.0520 

0.0407 
0.0460 
0.0423 
0.0455 
0.0402 

0.016 
0.024 
0.047 
0.076 
0.142 

0.024 
0.040 
0.066 
0.115 
0.167 

101 

1? 

75 

T> 

V 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

7)15° 

i)25° 

7)35° 

*15° 

«,25° 

^35° 

15  to  25° 

25  to  35° 

4 
10 
Solv. 

0.04688 
0.04496 
0.04301 

0.03527 
0.03376 
0.03221 

21.33 

22.25 
23.25 

28.35 
29.62 
31.05 

:::::::: 

0.02510 

39.84 

TABLE  27. — Cobalt  Bromide  in  Formamid. 


Molecular  conductivity 

Dissociation. 

Per  cent. 

Conductivity 
units. 

15° 

25° 

35° 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

15  to  25° 

25  to  35" 

10 

27.11 

35.35 

44.30 

60.9 

61.4 

60.7 

0.0304 

0.0253 

0.825 

0.895 

50 

34.70 

45.01 

56.56 

78.0 

78.1 

77.5 

0.0297 

0.0257 

1.031 

1.155 

200 

39.00 

50.34 

87.7 

87.4 

0.0291 

1.134 

800 

44.48 

57.60 

72.95 

100.0 

100.0 

100.0 

0.0295 

0.0266 

1.312 

1.535 

1,600 

41  83 

53  96 

68  70 

0  029O 

0  0273 

1  213 

1  474 

Viscosity  and  fluidity. 

Temperature  coeffi- 
cients of  fluidity. 

7715° 

7)25° 

7)  35°           v  15° 

*>25° 

^35° 

15  to  25° 

25  to  35° 

10 

0.04892 

0.03699 

0.02867        20.44 

27.03 

34.88 

0.0322 

0.0290 

Solv. 

0.04301 

0.03221 

0.02510       23.25 

31.05 

39.84 

0.0335 

0.0283 

In  this  laboratory,  during  the  past  fifteen  years,  solutions  in  the  fol- 
lowing pure  solvents  have  been  investigated :  Water,  methyl  and  ethyl 
alcohols,  acetone,  and  glycerol.  Among  the  results  of  this  work  are  the 
discovery  or  conformation  of  many  relations  between  the  conductivity 
and  viscosity  of  solutions,  the  dissociation  of  the  solutes,  and  the  die- 


Conductivities  and  Viscosities  in  Formamid. 


35 


lectric  constant,  association  factors,  and  viscosities  of  the  solvents.1 
Reference  to  table  7  (p.  24)  shows  that  formamid  is  a  solvent  markedly 
different  from  any  other  pure  solvent  investigated  in  regard  to  the  three 
constants  mentioned.  The  evidence  obtained  supports  and  confirms 
all  the  relations  and  conclusions  that  have  been  established  by  Jones 
and  his  co-workers,  from  their  investigations  with  other  pure  solvents. 
These  relations  will  be  discussed  in  turn. 

The  Thomson2-Nernst3  theory  is  that  the  forces  which  hold  the  atoms 
together  are  electrical;  hence  the  solvent  having  the  highest  dielectric 
constant  has  the  greatest  dissociating  power.  The  dissociating  power 
of  solvents  is  shown  in  two  ways:  first,  by  comparing  the  percentage 
dissociation  of  solutions  having  the  same  normality ;  second,  by  com- 
paring the  dilutions  at  which  complete  dissociation  is  reached.  Com- 
pare table  12  with  table  28. 

TABLE  28.— Potassium  Iodide  in  Water. 


Temperature 

V 

V-v 

a 

coefficient  per 

cent  at  25°-35°. 

2 

112.8 

76.6 

1.85 

8 

120.7 

82.0 

1.97 

2,048 

147.2 

100.0 

2.04 

Table  12  shows  that  the  dissociation  of  the  N/2  solution  of  potassium 
iodide  in  formamid  is  only  71  per  cent  at  25°.  From  the  viscosity  data 
in  table  12,  we  calculate  that  the  viscosity  of  the  N/2  solution  is  12.4 
per  cent  greater  than  the  solvent,  while  the  viscosity  of  the  N/2  solution 
of  potassium  iodide  in  water  is  known  to  be  less  than  that  of  the  solvent. 
From  a  large  mass  of  evidence,  we  know  that  viscosity  is  by  far  the 
largest  factor  affecting  conductivity  in  solutions  in  which  dissociation  is 
of  the  same  order  of  magnitude.  The  conductivity  of  the  N  /2  solution  in 
formamid  at  25°  is  18.67.  We  can  assume,  without  appreciable  error, 
that  if  the  viscosity  of  the  solvent  and  the  N/2  solution  were  in  the  same 
ratio  as  in  the  case  of  the  potassium  iodide  solution  in  water,  the  conduc- 
tivity of  the  N/2  solution  in  formamid  would  be  at  least  12.4  per  cent 
larger  than  the  figure  given — 18.67.  Recalculating  the  dissociation  of 
the  potassium  iodide-formamid  solution  on  this  basis  gives  79.4  per  cent, 
compared  with  76.6  per  cent  for  the  potassium  iodide  solution  in  water. 
This  latter  figure  would  be  even  less,  if  corrections  were  made  for  the 
fact  that  the  viscosity  of  an  N/2  solution  in  water  is  less  than  that  of 
the  solvent,  which  would  be  a  legitimate  correction  to  make  for  the 
purpose  of  this  comparison.  Comparing  in  this  way  the  dissociation 
of  these  two  N/2  solutions,  we  find,  as  would  be  expected  from  the 
Thomson-Nernst  theory,  that  formamid  has  the  greater  dissociating  power. 

'These  results  have  been  tabulated  and  discussed  in  publications  of  the  Carnegie  Institution  of 
Washington,  Nos.  170,  180,  and  210. 

'Phil.  Mas.,  36,  320  (1893).  3Zeit.  phys.  Chem.,  13,  531  (1894). 


36  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

Complete  dissociation  of  potassium  iodide  in  formamid  is  found 
in  the  N/200  solution,  while  in  water  it  is  not  reached  till  the  N/2048 
solution.  Therefore,  in  both  respects  formamid  falls  in  line  with  the 
Thomson-Nernst  theory.  This  relation  is  confirmed  by  every  salt 
studied  in  this  investigation  for  which  data  for  comparison  are  available. 
The  rubidium  salts  show  the  stronger  dissociating  power  of  formamid 
much  more  than  some  others  that  were  investigated. 

The  hypothesis  of  Dutoit  and  Aston1  states  that  the  greater  the  asso- 
ciation factor  of  the  solvent  the  greater  its  dissociating  powers.  Table 
7  shows  the  relation  of  formamid  to  other  pure  solvents  in  regard  to 
association  factors.  The  correction  and  comparisons  made  above, 
with  reference  to  the  Thomson-Nernst  theory,  are  equally  applicable 
in  connection  with  the  Dutoit  and  Aston  theory. 

Jones  and  Mahin  showed  that  complete  dissociation  of  lithium 
nitrate  in  acetone  is  not  reached  even  when  V  is  100,000.  The  signifi- 
cance of  these  figures  is  apparent,  if  one  has  in  mind  the  association 
factors  and  dielectric  constants  of  formamid,  ethyl  alcohol,  and  acetone, 
given  in  table  7. 

Jones  and  Veazey's2  explanation  of  the  decrease  in  viscosity  produced 
to  the  greatest  extent  by  caesium  and  rubidium  salts,  and  in  a  less 
degree  by  potassium,  ammonium,  and  other  salts,  all  having  very  large 
molecular  and  atomic  volumes,  has  been  of  special  interest  in  this 
investigation.  They  base  their  hypothesis  on  the  theory  of  Thorpe 
and  Rodger,  that  viscosity  is  due  to  the  friction  between  the  surfaces  of 
the  molecules.  If  the  particles  of  the  solute  are  larger  than  those  of  the 
solvent,  the  frictional  surfaces  and,  consequently,  the  viscosity  will  be 
decreased.  If  the  added  particles  are  smaller  than  those  of  the  solvent, 
the  viscosity  will  be  increased  by  the  addition  of  the  solute.  The  vis- 
cosity of  all  the  concentrated  solutions  of  salts  having  very  large 
atomic  or  molecular  volumes,  has  been  found  to  be  less  than  the  solvent 
in  the  case  of  all  pure  solvents,  except  acetone,  which  have  previously 
been  studied.  The  viscosity  of  formamid  has,  on  the  contrary,  been 
increased  by  every  salt  used.  The  complex  formamid  molecules 
HCONH2,  and  its  very  large  association  factor  (6.18)  show  that  its 
actual  molecule  is  larger  than  that  of  any  other  pure  solvent  used  in  these 
investigations,  and  the  increase  in  viscosity  by  the  salts  named  above 
confirms  this  conclusion.  The  following  comparison  strikingly  illus- 
trates this  relation.  A  N/2  solution  of  rubidium  iodide  in  glycerol 
decreases  the  viscosity  of  the  solvent  13  per  cent,  while  a  N/2  solution  of 
the  same  salt  in  formamid  increases  the  viscosity  of  the  solvent  12 
per  cent.  This  relation  will  be  referred  to  again  in  discussing  the 
results  of  the  work  with  caesium  nitrate  and  chloride. 

'Compt.  Rend.,  125,  240  (1897). 

2Amer.  Chem.  Journ.,  37,  405  (1907);  Carnegie  Inst.  Wash.  Pub.  No.  80  (1907). 


Conductivities  and  Viscosities  in  Farmamid.  37 

That  the  percentage  temperature  coefficients  of  conductivity 
increase  as  the  viscosity  of  the  solution  increases  is  a  relation  that  has 
been  brought  out  by  all  the  investigations  in  this  laboratory.  These 
coefficients  for  solutions  in  formamid  have  the  order  of  magnitude 
that  would  be  expected  from  the  relation  of  their  viscosities  to  the 
viscosities  of  the  same  solutions  in  other  solvents.  For  example, 
compare  the  coefficients  for  solutions  of  potassium  iodide  in  water,  table 
28,  with  those  for  the  same  salt  in  formamid  (table  12). 

All  salts  containing  water  of  crystallization  were  carefully  dehydrated 
at  suitable  temperatures  just  before  the  solutions  were  prepared. 
No  thermal  measurements  were  made  of  the  heat  of  solution,  but  in 
the  case  of  sodium  iodide,  which  crystallizes  with  two  molecules  of 
water,  a  marked  rise  in  temperature  was  noted  when  the  salt  dis- 
solved in  formamid,  which  indicated  the  formation  of  a  solvate.  It  is 
well  known  that  solvated  salts  give  higher  temperature  percentage 
coefficients  of  conductivity  than  non-solvated,  because  complexes  are 
usually  simplified  by  rise  in  temperature.  In  the  present  investigation, 
salts  crystallizing  with  water  have  given  the  larger  coefficients,  indi- 
cating the  formation  of  solvates  analogous  to  the  formation  of  solvates 
by  the  same  salts  in  water  and  other  solvents.  To  illustrate  this 
relation,  let  us  compare  these  coefficients  for  sodium  chromate  crystal- 
lizing with  10  molecules  of  water  (table  10)  and  cobalt  bromide  crystal- 
lizing with  2  molecules  of  water  (table  27).  Also,  compare  sodium 
iodide,  crystallizing  with  2  molecules  of  water  (table  9)  with  potassium 
iodide  (table  12).  The  coefficients  for  the  dilute  solutions  of  sodium 
chromate  are  probably  not  reliable  for  this  comparison,  as  will  be 
explained  under  the  discussion  of  this  salt. 

SODIUM  CHROMATE. 

Table  10  gives  the  results  for  sodium  chromate.  They  are  of  the 
usual  order  of  magnitude,  except  for  the  N/1600  solution.  At  each 
temperature  the  increase  in  conductivity  between  the  N/800  and 
N/1600  solutions  is  59  per  cent,  which  is  probably  due  to  chemical 
action  or  decomposition  instead  of  to  an  increase  in  ionization.  Formic 
acid  is  a  strong  reducing  agent.  Formamid  may  be  considered  as 
formic  acid  in  which  a  hydroxyl  has  been  replaced  by  an  amino  group. 
Both  the  hydroxyl  and  amino  group  are  basic;  therefore,  we  might 
expect  from  this  relationship  that  formamid  would  be  a  reducing 
agent.  The  reducing  action  of  formamid  will  also  be  discussed  with 
the  results  for  mercuric  chloride.  Sodium  chromate  is  a  strong  oxi- 
dizing agent;  hence,  it  is  not  surprising  that  there  should  be  chemical 
action  between  these  two  compounds  when  the  chromate  is  in  a  highly 
ionized  condition,  as  in  a  dilute  solution.  The  percentage  dissociations 
have  not  been  calculated,  because  the  conditions  just  referred  to  intro- 
duce an  uncertainty  regarding  them. 


38  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

MERCURIC  CHLORIDE. 

The  conductivity  of  aqueous  solutions  of  mercuric  chloride  is  too 
small  to  be  measured,  but  in  formamid,  on  account  of  its  greater  dis- 
sociating power,  the  conductivity  is  measurable  and  increases  as  dilu- 
tion increases  in  the  usual  way  up  to  the  N/400  solution.  The  con- 
ductivity values  for  the  N/1000,  N/1600  and  N  3200  solutions  clearly 
result  from  some  decomposition,  probably  of  the  solvent.  The  con- 
ductivity of  barium  chloride  at  complete  dissociation  is  55.86  at  25°. 
The  large  value  134.49  for  the  N/1000  solution  of  mercuric  chloride 
shows  that  it  is  not  due  to  dissociation.  Assuming  that  at  complete 
dissociation  the  conductivity  of  mercuric  chloride  would  be  of  the  same 
order  of  magnitude  as  of  barium  chloride,  the  dissociation  of  the  N/4 
solution  would  be  about  1  per  cent,  increasing  to  about  7  per  cent  for 
the  N/400  solution.  Silver  chloride  is  precipitated  in  the  N/400  solu- 
tion by  silver  nitrate.  By  standing  in  the  sunlight  for  a  short  time,  a 
heavy  precipitate  of  metallic  silver  is  formed,  showing  again  the  reduc- 
ing action  of  formamid. 

To  recover  the  solvent  from  the  mercuric  chloride  by  vacuum  dis- 
tillation required  several  more  distillations  than  usual.  Metallic 
mercury  was  found  in  the  receiver  after  the  first  distillation,  showing 
again  the  reducing  action  of  formamid.  It  is  possible  that  the  liber- 
ated chlorine  formed  a  salt  with  the  formamid  which  was  easily  dis- 
sociated, since  the  conductivity  of  the  first  fraction  of  the  distillate 
was  very  high  and  was  lowered  only  a  small  amount  by  successive  dis- 
tillations. Repeated  distillation,  however,  yielded  the  solvent  used 
for  the  caesium  solutions. 

COBALT  BROMIDE. 

The  green,  anhydrous  cobalt  bromide  gives  a  pink  solution  in  forma- 
mid as  in  water.  After  the  determinations  were  made,  a  mixture  of 
solutions  was  distilled  to  recover  the  formamid.  The  mixture  con- 
tained mercuric  chloride  and  cobalt  bromide.  As  already  stated,  the 
mercuric  chloride  was  reduced  to  metallic  mercury.  The  dry  residue 
in  the  distillation  flask  had  the  brilliant  blue  color  of  anhydrous  cobalt 
chloride,  showing  that  the  free  chlorine  from  the  mercuric  chloride  had 
replaced  the  bromine  in  the  cobalt  bromide. 

CESIUM  NITRATE  AND  CHLORIDE. 

For  several  years  past  this  laboratory  has  been  unable  to  obtain  any 
caesium  salts  for  its  investigations  on  conductivity  and  viscosity. 
Recently  some  caesium  sulphate  was  obtained  through  the  cooperation 
and  courtesy  of  Prof.  James  Lewis  Howe,  of  Washington  and  Lee  Uni- 
versity. The  sulphate  was  converted  into  the  nitrate  and  chloride, 
and  the  conductivity  and  viscosity  of  solutions  of  these  salts  were 
determined.  The  results  are  recorded  in  tables  22  and  23. 


Conductivities  and  Viscosities  in  Formamid.  39 

A  survey  of  the  work  on  the  salts  of  all  the  alkali  metals  shows 
that  conductivity  values  for  the  caesium,  rubidium,  and  potassium  salts 
are  approximately  the  same,  and  for  sodium  and  lithium  are  less. 
The  same  relation  is  true  for  the  results  obtained  in  aqueous  solutions. 
Dissociation  percentage  is  of  the  same  order  of  magnitude  for  all  the 
salts  named.  The  conductivity  results,  in  connection  with  the  disso- 
ciation percentage,  harmonize  with  the  ionic  velocities.  For  caesium, 
rubidium,  and  potassium  the  relative  ionic  velocities  are  73.6,  73.5, 
and  70.6,  and  for  sodium  and  lithium  they  are  49.2  and  39.8. 

The  most  striking  difference  between  these  metals  is  shown  by  the 
atomic  volume  curve  of  Lothar  Meyer.1  This  function  increases  by 
nearly  regular  steps  from  12  for  lithium  to  72  for  caesium.  In  accord- 
ance with  the  theory  of  Jones  and  Veazey,2  to  which  reference  has  already 
been  made,  the  percentage  increases  in  the  viscosity  of  the  concentrated 
solutions  over  the  viscosity  of  the  solvent,  should  be  the  smallest  for 
caesium  salts,  and  should  increase  in  the  following  order;  caesium,  rubid- 
ium, potassium,  sodium,  and  lithium.  This  relation  comes  out  clearly  by 
comparing  the  figures  given  in  the  last  column  of  tables  8  to  27.  (Com- 
pare these  results  for  caesium  and  rubidium  salts  with  those  for  sodium 
and  lithium) .  The  salts  of  rubidium  and  caesium  increase  the  viscosity 
of  formamid  much  less  than  sodium  and  lithium  salts.  The  difference 
in  percentages  of  increase  is  not  so  marked  as  the  difference  in  the 
atomic  volumes  of  the  metal  ions.  The  effect  is  partially  suppressed 
by  other  factors.  Consider  the  solution  of  a  binary  salt  which  is  80 
per  cent  dissociated.  Every  hundred  molecules  of  the  salt  gives,  in 
solution,  80  anions,  80  cations,  and  20  molecules.  The  80  cations  are 
only  44.5  per  cent  of  the  number  of  particles  in  the  solution.  It  is 
only  this  44.5  per  cent  of  cations  which  have  the  atomic  volume  rela- 
tions referred  to  above.  The  molecular  volumes  do  not  show  this 
relation,  and  the  atomic  volumes  of  the  anions  is  unknown. 

SUMMARY  OF  RESULTS. 

The  first  six  of  the  conclusions  drawn  below  have  been  developed 
or  confirmed  by  the  investigations  in  this  laboratory  with  other  pure 
solvents,  and  have  been  shown  by  this  investigation  to  hold  true  for 
formamid. 

1.  The  greater  the  dielectric  constant  of  a  solvent  the  greater  its 
dissociating  power. 

2.  The  greater  the  association  factor  of  a  solvent  the  greater  its 
dissociating  power. 

3.  The  formation  of  solvates  with  formamid  is  indicated  by  those  salts 
that  form  hydrates  with  water. 

'Lieb.  Ann.  Suppl.,  7,  354  (1870). 

'Carnegie  Inst.  Wash.  Pub.  No.  80  (1907);  Amer.  Chem.  Journ.,  37,  405  (1907). 


40  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

4.  Solvated  salts  show  larger  percentage  temperature  coefficients  of 
conductivity  than  non-solvated  salts. 

5.  The  order  of  magnitude  of  percentage  temperature  coefficients 
of  conductivity  is  approximately  proportional  to  the  viscosity  of  the 
solutions. 

6.  Formamid  is  the  first  of  the  pure  solvents  studied  with  reference 
to  the  viscosity  of  solutions,  in  which  none  of  the  salts  used  produces 
negative  viscosity.     A  satisfactory  explanation  of  the  above  fact  is 
apparent,  if  the  very  large  association  factor  of  the  solvent  is  considered 
in  connection  with  the  theory  of  Jones  and  Veazey. 

7.  In  regard  to  conductivity,  dissociation  and  effect  on  viscosity 
of  solvent,  caesium  salts  are  very  closely  allied  to  rubidium  salts. 

8.  Mercuric  chloride  is  more  dissociated  in  formamid  than  in  water. 


CHAPTER  III. 

RADIOMETRIC  MEASUREMENTS  OF  THE  IONIZATION  CONSTANTS  OF 

INDICATORS.-I. 


BY  E.  J.  SHAEFFER  AND  M.  G.  PAULUS. 


The  radiometric  measurements  recorded  in  this  paper  were  made 
with  a  very  sensitive  radiomicrometer  and  a  grating  spectroscope. 
The  spectroscope  is  well  known  to  the  chemist,  but  its  use  in  connection 
with  the  radiomicrometer  has  not,  as  yet,  found  an  extensive  applica- 
tion to  the  study  of  chemical  problems.  This  is  perhaps  due  in  part  to 
the  difficulty  in  constructing  a  sensitive  and  easily  controlled  radio- 
micrometer.  By  means  of  the  grating  spectroscope  and  radiomicrom- 
eter, it  is  possible  to  study  quantitatively  chemical  reactions  involving 
color  changes,  and  even  those  not  involving  such  changes,  if  there  are 
absorption  bands  in  the  invisible  regions  of  the  spectrum.  It  will 
be  seen  that  accurate  determinations  of  very  small  concentrations  of 
colored  components  in  solutions  can  be  made  very  rapidly  even  when 
two  or  more  such  components  are  present.  The  structure  of  the  solvent 
bands  and  time  reactions  are  among  the  many  other  important  chem- 
ical problems  which  may  be  investigated  by  means  of  this  radiometric 
apparatus. 

PURPOSE  OF  THIS  INVESTIGATION. 

In  connection  with  some  problems  now  under  investigation  in  this 
laboratory,  it  was  desired  to  secure  accurate  measurements  of  the 
hydrolysis  constants  of  certain  salts.  The  spectroscope  and  radio- 
micrometer  suggested  the  indicator  method,  which  is  based  on  the 
changes  in  the  transmission  of  light  by  solutions  of  indicators  when 
varying  amounts  of  hydrogen  orhydroxyl  ions  are  present.  It  will  be 
shown  that  by  means  of  these  changes  the  concentrations  of  very  small 
amounts  of  hydrogen  and  hydroxyl  ions  can  be  determined.  The 
method  obviously  involves  a  knowledge  of  the  ionization  constant  of 
the  indicator  as  an  acid  or  as  a  base.  A  comparison  of  the  ionization 
constants  of  the  indicators  obtained  by  several  investigators,  showed 
that  the  recorded  values  varied  widely,  not  only  with  different  methods, 
but  with  the  same  method.  The  results  for  phenolphthalein,  in  partic- 
ular, indicate  how  inaccurate  is  our  knowledge  as  to  the  ionization 
constant  of  that  important  indicator.  Salm1  gives  the  value  8X10~10; 
McCoy,2  0.8X10"10.  In  many  cases  the  results  of  an  individual 
investigation  vary  as  much  as  300  to  400  per  cent.  A.  A.  Noyes3 

'Zeit.  phys.  Chem.,  57,  492  (1907).  Mourn.  Amer.  Chem.  Soc.,  32, 858  (1910). 

"Arner.  Chem.  Journ.,  31,  503  (1904). 

41 


42  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

has  applied  the  theory  of  indicators  to  volumetric  analysis,  making  use 
of  the  ionization  constants  of  the  various  indicators.  He  expresses 
the  opinion  that  our  knowledge  of  the  ionization  constants  of  the 
indicators  is  for  the  most  part  inexact  and  needs  to  be  supplemented 
by  further  careful  investigation.  Stieglitz,1  McCoy,2  and  Salm  have 
also  recognized  the  importance  of  such  a  study.  McCoy,  Salm,  and 
in  particular  Noyes  and  Bjerrum,3  have  already  applied  the  indicator 
constants  to  some  of  the  more  important  problems  which  the  chemist 
must  face  in  titrating  weak  acids  and  bases;  and  they  have  shown  in 
certain  cases  the  extent  of  the  error  involved  in  such  titrations.  It  will 
then  be  readily  seen  that  it  is  important  to  know  the  ionization  con- 
stants of  the  more  important  indicators,  especially  of  methyl  orange 
and  phenolphthalein.  In  view  of  these  facts,  and  the  value  of  such 
ionization  constants  in  many  other  lines  of  work  involving  the  use  of 
indicators,  it  was  decided  to  make  a  radiometric  investigation  of  this 
problem.  Methyl  orange  was  first  studied,  and  a  method  was  devel- 
oped through  the  application  of  Beer's  law,  which  readily  gave  the 
concentration  of  the  yellow  azo-base  and  that  of  the  red  quinoid  ion. 
Knowing  these  values  and  the  hydrogen  ion  concentration,  the  hydroly- 
sis constant  of  methyl  orange  can  be  readily  calculated.  The  method 
employed  is  different  from,  and  was  worked  out  independently  of  any 
other  method  thus  far  used  in  investigating  this  problem. 

HISTORICAL. 

The  results  of  a  number  of  preliminary  investigations  on  the  ioni- 
zation constants  of  various  indicators  were  published  in  1904.  McCoy, 
like  most  of  the  other  investigators  who  studied  this  problem,  worked 
calorimetrically,  the  intensity  of  the  color  being  judged  by  the  eye. 
McCoy  employed  Nessler  tubes  for  this  purpose,  while  others  made  use 
of  special  calorimeters  to  determine  the  amount  of  indicator  trans- 
formed into  its  salt  by  varying  amounts  of  hydrogen  or  hydroxyl  ions. 
Tables  showing  roughly  the  hydrogen  ion  concentration  at  which  a 
large  number  of  indicators  undergo  change  in  color,  have  been  pub- 
lished by  Friedenthal,4  Salessky,5  Fels,6  and  Salm.7  Nearly  all  of  the 
early  investigators  utilized  the  principle  that  the  ionization  constant 
of  an  indicator  is  equal  to  the  hydrogen  ion  concentration  at  which  it 
is  one-half  transformed  into  its  salt.  (See  Salm  and  A.  A.  Noyes.) 
Salessky,  and  later  Salm,  determined  the  concentration  of  the  hydrogen 
ions  at  this  point  by  means  of  the  hydrogen  electrode.  Criticisms 
showing  how  inexact  most  of  these  methods  are  will  be  found  in  the 
separate  articles.  Salm,8  referring  to  the  work  previous  to  1907,  says 

'Journ.  Amer.  Chem.  Soc.,  25,  1126  (1903).  *Ibid.,  10,  204  (1904). 

2Amer.  Chem.  Journ.,  31,  503  (1904).  'Ibid.,  10,  208  (1904). 

•Ahr.  Versamm.,  21,  1  (1914).  '/bid.,  10,  344  (1904). 

4Zeit.  Elektroohem.,  10,  113  (1904).  »Zeit.  phys.  Chem.,  57,  490  (1907). 


PLATE  2 


1 .     Spectroscope  without  Cover. 


2.     Radiomicrometer. 


Radiometric  Measurements  of  Constants  of  Indicators.  43 

that  the  investigations  have  been  for  the  most  part  qualitative,  and  that 
with  few  exceptions  the  dissociation  constants  of  the  indicators  are 
still  unknown. 

Salm  made  a  more  careful  study,  employing  a  satisfactory  calori- 
meter to  determine  when  the  indicator  was  one-half  transformed  into 
its  salt,  and  like  Salessky  obtained  the  hydrogen  ion  concentration 
at  this  point  by  means  of  the  hydrogen  electrode.  His  results  with 
phenolphthalein  illustrate  the  uncertainty  of  the  values  found  by  his 
method.  Wegscheider1  proceeded  in  much  the  same  manner  as  the 
others  who  have  used  the  colorimetric  method,  and  obtained  con- 
stants for  the  ionization  of  phenolphthalein  which  seem  to  be  in  fair 
agreement  with  those  found  by  Hildebrand.2  Hildebrand's  method, 
involving  the  use  of  the  spectro-photometer  for  the  estimation  of  color 
intensities,  is  the  most  exact  of  all  the  methods  previously  employed. 
His  photometric  method  somewhat  resembles  the  radiometric  method 
which  we  employed.  Rosenstein,3  following  essentially  the  procedure 
of  McCoy,  has  carried  out  very  carefully  a  colorimetric  investigation  of 
the  ionization  constant  of  phenolphthalein,  and  the  effect  upon  it  of 
neutral  salts.  He  employed  the  Duboscq  type  of  colorimeter  to  deter- 
mine the  fraction  of  the  indicator  transformed  by  a  known  hydrogen 
ion  concentration.  As  he  shows,  the  ionization  constant  is  equal  to 
the  hydrogen  ion  concentration  multiplied  by  the  fraction  of  the  indi- 
cator transformed  into  its  salt.  When  the  indicator  is  a  fairly  strong 
electrolyte,  e.  g.,  p-nitrophenol,  its  dissociation  constant  was  deter- 
mined by  the  conductivity  method. 

Before  we  can  judge  of  the  absolute  value  of  the  results  obtained  in 
any  radiometric  investigation,  it  is  necessary  that  the  radiometric 
instrument  and  the  other  parts  of  the  apparatus  used  in  connection 
with  it,  which  fulfill  several  requirements,  be  taken  up  in  the  discus- 
sion of  the  apparatus  which  follows.  In  view  of  the  fact  that  radio- 
metric  apparatus  has  not  thus  far  been  extensively  used  by  chemists, 
it  seems  desirable  that  this  be  discussed  in  some  detail. 

Photographs  of  the  assembled  apparatus  and  also  of  the  various 
parts  are  given  in  plate  1,  figure  2,  and  plates  2  and  3  show  the  assem- 
bled apparatus.  The  long  box  to  the  left  (plate  1,  fig.  2)  incloses  the 
spectroscope.  The  standard  which  holds  the  Nernst  glower,  lenses, 
sliding  carriage,  etc.,  occupies  the  center,  and  the  tall  box  to  the  right 
incloses  the  radiomicrometer.  Below  the  standard  is  shown  the  rheostat 
ammeter,  etc.,  for  controlling  the  current  supplied  to  the  Nernst  glower. 
Plate  2,  figure  1,  is  a  photograph  of  the  apparatus  with  the  covers 
removed  from  the  spectroscope  and  the  radiomicrometer.  To  the 
extreme  left  is  shown  the  grating  and  a  4-inch  Brashear  lens,  and  at 
the  extreme  right  is  the  tube  containing  the  radiomicrometer  with  the 

'Zeit.  Elektrochem.,  14,  510  (1908).  'Ibid.,  14,  351  (1908). 

Mourn.  Amer.  Chem.  Soe.,  34,  1117  (1912). 


44  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

permanent  magnet  at  the  bottom.  Plate  2,  figure  2,  is  a  side  view  of 
the  radiomicrometer.  A  larger  view  of  the  lens  and  grating  is  shown 
in  plate  3,  figure  1 ;  an  end  view  of  the  spectroscope  in  plate  3,  figure  2. 

THE  RADIOMICROMETER. 

The  methods  employed  in  the  construction  of  the  radiomicrometer 
and  especially  of  the  thermo-j  unctions,  we  owe  to  Professor  A.  H.  Pfund.1 
By  means  of  these  methods  we  were  able  to  build  a  very  satisfactory- 
instrument. 

The  radiomicrometer  previously  constructed  by  Guy2  was  found  to 
be  unsatisfactory  for  use  with  the  grating  spectroscope.  Its  sensibility, 
according  to  a  recent  test,  was  2  per  square  millimeter  of  exposed  vane, 
candle  and  scale  being  at  a  meter's  distance.  The  full  period  of  this 
instrument  was,  however,  very  short,  being  only  8  seconds.  His  radio- 
micrometer  could  have  been  made  about  as  sensitive  as  the  one  con- 
structed for  this  work,  by  using  a  longer  and  finer  quartz  fiber  and 
making  the  full  period  about  20  seconds.  However,  as  this  instrument 
was  not  equipped  with  a  compensating  junction,  and  as  its  drift  due  to 
this  cause,  had  proved  so  troublesome  in  the  work  of  the  previous  year, 
it  was  decided  to  construct  a  new  radiomicrometer. 

Briefly,  the  radiomicrometer  consists  of  a  thermo-electric  junction 
attached  to  a  loop  of  non-magnetic  wire.  The  whole  system  is  sus- 
pended by  a  quartz  fiber  in  a  glass  tube.  A  strong  magnetic  field 
surrounds  that  portion  of  the  tube  inclosing  the  loop  of  wire.  Radiant 
energy  falling  upon  the  blackened  junction  is  converted  into  electrical 
energy.  In  proportion  to  the  amount  of  energy  received  by  the 
junction,  the  suspended  system  turns  through  a  definite  angle  in  the 
magnetic  field,  the  loop  tending  to  set  itself  at  right  angles  to  the  lines 
of  force.  The  deflection  or  turn  is  given  by  means  of  a  mirror  attached 
to  the  suspended  system. 

The  essential  parts  of  the  radiomicrometer  are  shown  in  figure  6, 
which  represents  the  type  of  compensating  junction  that  was  con- 
structed and  used.  It  is  fastened  to  the  loop  of  copper  wire  w  at  c 
and  c',  as  shown  in  figure  6.  To  support  the  loop  of  wire  small  glass 
rods  are  placed  at  h  and  ti.  A  light  mirror  is  attached  to  h  at  m. 
A  quartz  fiber,  /,  is  fastened  to  the  end  of  the  glass  rod  h,  and  to  the 
brass  stopper  S.  The  whole  system  is  then  set  in  the  glass  tube  T, 
having  two  windows  I  and  p.  The  lens  I  focuses  the  reflected  light 
from  the  mirror  m,  on  a  glass  scale.  At  p  is  inserted  a  plane-glass 
window  1  mm.  thick.  The  beam  of  light  which  is  to  be  measured 
passes  through  this  window  before  falling  on  the  junction.  The  mag- 
netic field  is  placed  between  I  and  p. 

'Phys.  Rev.,  34,  228  (1912);  Phys.  Zeit.,  13,  870  (1912). 
'Carnegie  Inst.  Wash.  Pub.  No.  190,  30  (1913). 


PLATE  3 


I .     Lens  and  Grating  of  Spectroscope. 


2.     Spectroscope  end  View. 


Radiometric  Measurements  of  Constants  of  Indicators. 


45 


The  wire  loop  was  made  of  a  very  fine  specimen  of  No.  36  copper 
wire,  furnished  by  Leeds  and  Northrup.  Pure  nitric  acid  was  used  to 
dissolve  the  surface  of  the  wire,  which  was  very  likely  to  be  con- 
taminated with  various  magnetic  materials.  That  this  wire  must 
have  been  very  pure  is  shown  by  the  behavior  of  the  completed  instru- 
ment. As  regards  purity,  a  still  better  specimen  of  silver  wire  was 
obtained  from  Weston;  but  unfortunately  none  of  the  junctions 
attached  to  this  silver  wire  was  sufficiently  sensitive  for  our  purpose. 

The  alloys  used  for  the  construction  of  the  thermo-j  unction  were  of 
the  composition  recommended  by  Hutchins.1  Figure  6  of  the  sketch 
gives  an  enlarged  view  of  the  junction.  It  will  be  noticed  that  it 
consists  of  two  thermo-electric  junctions  A  and  B,  which  compensate 
each  other.  When  the  same  beam  of  light  falls  on  both  junctions,  each 
sends  equal  amounts  of  electrical  energy 
through  the  loop  of  non-magnetic  wire  w, 
but  in  opposite  directions.  Hence,  it  follows 
that  for  exact  compensation  there  should  be 
no  turn  of  the  suspended  system.  In  actual 
operation  the  beam  of  light  energy  which  is 
to  be  measured  is  focused  on  one  of  the 
junctions.  This  arrangement  of  junctions 
is  quite  essential  in  any  sensitive  radio- 
micrometer  where  a  constant  zero-point  is 
required;  and  it  is  only  in  this  way  that 
the  effects  of  temperature  changes  and  stray 
sources  of  light  can  be  eliminated. 

The  two  arms  represented  by  a  and  a' 
in  figure  6  have  the  composition  97  parts 
bismuth  and  3  parts  antimony.  The  com- 
position of  6  is  95  parts  bismuth  and  5  parts 
tin.  These  metals  are  quite  pure  and  the 
strips  used  are  very  thin.  Junction  A  is  formed  by  sealing  strip  a  to 
strip  b,  and  junction  B  is  formed  by  sealing  strip  b  to  strip  a'.  R  and 
R'  are  the  receiving  surfaces,  or  so-called  vanes,  cut  from  very  thin 
tin  foil,  and  soldered  to  junctions  A  and  B,  respectively.  These 
vanes  each  have  an  area  of  4  sq.  mm.,  and  are  coated  with  lampblack 
to  prevent  radiation  of  light  energy.  It  was  found  that  the  action  of 
an  acidified  solution  of  antimony  chloride  on  the  tin  foil  produces  a 
black  receiving  surface.  Theoretically,  it  appears  that  such  a  metallic 
receiving  surface  should  be  more  effective  than  lampblack.  However, 
none  of  the  junctions  having  the  antimony  receiving  surface  was 
sufficiently  sensitive  to  determine  whether  or  not  its  use  is  to  be  pre- 
ferred to  that  of  lampblack.  Greater  sensibility  can  be  obtained  with 
a  single  junction  than  with  a  compensating  one,  since  in  the  former 


O" 


i 


I 


Fio.  6. 


•Sill.  Journ.,  48,  226  (1894). 


46  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

we  have  but  one  seal  and  two  metal  strips.  This  follows  from  the  fact 
that  the  sensibility  of  the  radiomicrometer  is  materially  increased 
when  the  resistance  of  the  junction  is  made  more  nearly  equal  to  that 
of  the  wire  loop.  However,  the  advantages  to  be  derived  from  the 
compensating  type  of  junction  are  well  worth  the  sacrifice  in  sensibility. 

It  is  very  important  that  the  weight  of  the  junction  be  kept  as  small 
as  possible,  not  only  to  lessen  the  weight  of  the  suspended  system,  but 
especially  to  reduce  the  heat  capacity  of  the  junction  to  a  minimum.  A 
small  heat  capacity  insures  a  quick-acting  suspension,  and  one  which 
will  more  rapidly  return  to  the  zero-point.  That  the  mass  of  material 
in  the  junction  is  very  small  can  be  seen  from  the  fact  that  the  actual 
weight  of  one  of  the  completed  compensating  junctions  was  2.9  mg. 

The  quartz  fiber1  was  obtained  from  molten  quartz  by  means  of  a  bow- 
and-arrow  arrangement.  The  total  length  of  this  fiber  is  nearly  30  cm. 
A  long  fiber  materially  aids  in  eliminating  vibrations.  This  is  most 
essential,  especially  when  the  work  necessitates  small  deflections.  The 
period  and  sensibility  of  the  radiomicrometer  depend  in  a  large  measure 
on  the  proper  selection  of  the  quartz  fiber. 

The  plane  mirror  m  (fig.  6)  is  very  thin  and  has  an  area  of  about  20 
sq.  mm.  Directly  in  front  of  it  is  the  lens  I,  having  a  focal  length  of  12 
feet.  By  means  of  this  arrangement  the  image  of  a  lamp  filament  could 
be  sharply  focused  on  a  glass  scale  about  12  feet  distant  from  the 
mirror;  and  it  is  on  this  scale  that  the  deflections  as  given  by  the  radio- 
micrometer  are  noted. 

The  torsion-head  stopper  s,  made  of  brass  and  tightly  ground  into  the 
glass  tube  T,  serves  the  purpose  of  making  adjustment.  By  means  of 
it  the  suspended  system  can  be  made  to  occupy  any  position  in  the 
magnetic  field,  and  the  whole  system  can  be  raised  or  lowered  without 
danger  of  breaking  the  very  fragile  suspension.  The  torsion-head 
stopper  contains  a  brass  rod  r  and  the  two  screws  st  and  sa  for  making 
the  above  adjustments. 

The  glass  tube  T,  in  which  the  whole  system  is  suspended,  is  about 
45  cm.  in  length.  To  it  is  attached  the  lens  I,  previously  referred  to, 
and  the  glass  window  p,  through  which  the  beam  of  light  is  directed 
either  to  junction  A  or  to  junction  B. 

The  glass  tube  surrounding  that  portion  of  the  suspended  system  to 
which  the  thermo-j  unction  is  attached  was  insulated  from  temperature 
changes  and  air  drafts  by  a  layer  of  fine  lead  shot  and  wool  fiber.  To 
prevent  stray  sources  of  light  and  air  drafts  from  reaching  the  radio- 
micrometer,  it  was  inclosed  in  a  wooden  box  covered  with  painted 
canvas. 

That  the  image  on  the  scale  should  remain  steady  at  a  focal  distance 
of  12  feet,  it  is  necessary  to  shield  the  radiomicrometer  very  carefully 
from  the  usual  vibrations  of  the  building.  The  long  quartz  fiber  aided 

'Dr.  C.  W.  Hewlett  baa  very  kindly  supplied  us  with  several  satisfactory  fibers. 


Radiometric  Measurements  of  Constants  of  Indicators.  47 

materially  in  this  respect.  Between  the  leveling  stand  of  the  radio- 
micrometer  and  the  base  on  which  it  rested,  various  insulating  mate- 
rials were  introduced  at  six  points  to  absorb  vibrations. 

The  completed  radiomicrometer,  candle  and  scale  being  at  a  meter's 
distance,  gave  a  deflection  of  20  cm.,  the  half-period  being  nearly  10 
seconds.  In  making  this  determination  of  the  sensibility  of  the  instru- 
ment the  radiation  was  passed  through  a  glass  window  1  mm.  thick, 
and  the  tube  containing  the  suspension  was  not  evacuated.  Evacuat- 
ing the  tube  and  passing  the  radiation  through  a  rock-salt  window 
would  increase  the  sensibility  about  six  times,  and  the  period  would 
become  somewhat  shorter. 

Since  the  receiving  vane  has  an  area  of  4  sq.  mm.,  the  sensibility  per 
square  millimeter  of  exposed  vane  is  5,  and  the  full  period  20  seconds. 
When  making  this  determination,  owing  to  the  difficulty  of  properly 
shielding  the  junction  not  in  use  but  serving  as  compensator,  it  is 
thought  that  the  value  of  5  for  the  sensibility  is  somewhat  low.  Prac- 
tically every  junction  that  was  tested  showed  that  the  sensibility  was 
approximately  one-fourth  of  the  whole  period  in  seconds.  The  sensi- 
bility value  as  given  by  the  candle  has  but  little  meaning,  owing  to  the 
varying  intensity  of  the  light  emitted.  The  light  from  the  Nernst 
glower,  burning  at  0.8  ampere  and  120  volts,  dispersed  into  the  visible 
first-order  spectrum  6  inches  long  at  slit  Si  of  the  spectroscope,  gave  at 
X  =  0.95/x  a  deflection  of  250  mm.  when  focused  on  the  junction  of  the 
radiomicrometer.  Both  slits  are  1  mm.  in  width,  and  the  scale  12  feet 
from  the  mirror.  Considering  the  high  resolving  power  of  the  grating, 
it  will  be  seen  that  our  radiomicrometer  is  very  satisfactory,  both  as 
regards  sensibility  and  period.  It  enabled  us  to  make  quantitative 
measurements  of  radiant  energy  from  X  =  0.4ju  to  X  =  2.0/i. 

For  the  same  source  of  light  energy,  the  radiomicrometer  always  gave 
the  same  deflection  and  returned  quickly  to  zero,  showing  that  our 
efforts  to  eliminate  magnetic  materials  from  the  wire  loop  had  been 
very  successful.  The  conduct  of  the  completed  instrument  has  demon- 
strated that  the  suspended  system  is  very  free  from  paramagnetic 
disturbances.  The  diamagnetic  materials  composing  the  junction  were 
all  arranged  at  right  angles  to  the  lines  of  force  of  the  magnetic  field, 
and  in  this  way  diamagnetic  disturbances  were  reduced  to  a  minimum. 
The  use  of  the  radiomicrometer  has  shown  that  the  heat  capacity  of 
the  junction  is  sufficiently  low  to  be  within  the  desired  limit.  The 
compensating  junction  and  its  insulation  from  temperature  changes 
practically  removed  any  drift  of  the  zero-point.  The  zero-point  for 
weeks  at  a  time  would  not  drift  more  than  10  cm.  on  either  side  of  the 
zero.  In  case  there  was  any  drift  during  the  measurements,  it  was 
always  very  slight  and  the  proper  correction  could  easily  be  applied. 
The  insulation  from  vibrations  made  the  reflected  beam  upon  the  scale 
fairly  steady,  even  when  there  were  rather  violent  disturbances  in  the 


48  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

building.  There  is  every  reason  to  suppose  that  the  scale  could  be 
read  to  about  0.25  mm.  when  the  building  was  quiet.  Duplicate 
readings  under  ordinary  conditions  nearly  always  agreed  to  within 
0.5  mm. 

THE  SPECTROSCOPE. 

The  grating  spectroscope  was  designed  and  built  in  the  Physical 
Laboratory  under  the  direction  of  Professor  John  A.  Anderson.  The 
use  of  such  a  spectroscope  presents  two  especially  desirable  advantages. 
First,  the  position  of  regions  of  absorption  can  be  very  accurately 
determined;  second,  due  to  its  high  resolving  power,  the  structure  of 
even  the  very  narrow  absorption  bands  and  some  absorption  lines  can 
be  studied  in  detail.  The  spectroscope  is  so  constructed  that  either  the 
photographic  plate  or  any  of  the  radiometric  instruments  can  be  used 
with  it.  The  use  of  the  radiomicrometer  enables  us  to  determine  not 
only  the  actual  positions  of  regions  of  absorption,  but  also  to  make 
quantitative  measurements  of  the  light  transmitted  by  a  solution  for 
a  wide  range  of  wave-lengths.  The  deflections  of  the  radiomicrometer 
give  an  accurate  measure  of  the  relative  intensities  of  the  different 
absorption  bands,  and  of  the  different  parts  of  the  same  bands.  It 
will  thus  be  seen  that  the  radiometric  method  has  distinct  advantages 
over  the  photographic,  which  is  chiefly  useful  in  determining  the 
positions  of  regions  of  absorption  and  the  general  characteristics  of  the 
visible  spectrum  as  transmitted  by  the  solutions.  Moreover,  the  photo- 
graphic plate  is  only  sensitive  from  X  =  0.2 /*  to  X  =  0.76 /i,  whereas,  with 
the  radiomicrometer  and  the  apparatus  used  in  this  work,  quantitative 
measurements  of  absorption  could  be  made  from  X  =  0.4;u  to  X  =  2.0ju, 
using  slits  only  1  mm.  in  width.  Considering  the  high  dispersion  that 
could  be  obtained  with  the  4-inch  grating,  this  width  of  slit  gives  a 
very  pure  spectrum. 

The  plane  4-inch  grating  with  which  the  spectroscope  is  equipped 
was  ruled  by  Anderson;  and  the  ruling  is  of  such  a  character  that  a 
very  bright  first-order  spectrum  is  produced.  This  is  quite  essential 
for  radiometric  measurements  in  the  visible  region  of  the  spectrum. 
Energy  measurements  were  made  in  both  first-order  spectra,  one  being 
on  each  side  of  the  central  image.  The  spectrum  on  one  side  was  found 
to  be  somewhat  more  intense,  and  therefore  this  brighter  side  was  used. 

If  either  the  intensity  of  the  light  source  or  the  sensitiveness  of  the 
radiomicrometer  exceeds  certain  limits,  it  is  possible  that  light  energy 
from  the  second-order  spectrum  will  vitiate  the  energy  measurements 
made  in  the  first-order  spectrum.  This  will  be  readily  seen  if  we  con- 
sider that  wave-length  of  light  X  =  0.35)u  of  the  second  order  overlaps 
wave-length  of  light  X  =  0.?M  of  the  first  order,  etc.  Therefore,  it 
was  desired  to  know  if,  with  a  Nernst  glower  burning  at  0.8  ampere 
and  120  volts,  the  second-order  spectrum  had  sufficient  energy  to  be 


Radiometric  Measurements  of  Constants  of  Indicators.  49 

detected  by  the  radioinicrometer.  If  such  proved  to  be  the  case,  it 
would  be  found  necessary  to  introduce  color  screens  to  eliminate 
the  light  energy  from  the  second-order  spectrum.  That  the  first- 
order  spectrum  can  be  regarded  as  pure  is  shown  by  the  following 
considerations: 

Glass  cuts  off  all  wave-lengths  of  light  beyond  X  =  0.35ju.  Since  the 
glass  in  the  path  of  the  light  is  of  considerable  thickness,  the  first-order 
spectrum  is  not  contaminated  with  any  light  from  the  second-order 
spectrum  as  far  out  as  X  =  0.7/t  of  the  first  order.  A  solution  of  copper 
sulphate  is  entirely  transparent  to  all  wave-lengths  of  light  shorter 
than  X  =  0.6//.  A  10  mm.  depth  of  a  saturated  solution  of  copper 
sulphate  shows  complete  absorption  beyond  X  =  0.6  n.  Using  the  above 
solution  of  copper  sulphate,  radiometric  measurements  were  made 
from  X  =  0.4)u  to  X  =  1.3/1-  Beyond  X  =  0.6/u  the  solution  of  copper 
sulphate  is  entirely  opaque,  and  therefore  no  evidence  of  light  energy 
from  the  second  order  was  obtained  as  far  out  as  X  =  1.3ju-  It  has 
already  been  shown  that  the  first-order  spectrum  is  pure  up  to  X  =  OJ/i, 
and  the  above  shows  that  the  light  from  the  overlapping  second-order 
spectrum  does  not  contain  a  sufficient  amount  of  energy  to  be  detected 
by  the  radiomicrometer  as  far  out  in  the  infra-red  as  X  =  1.2)u  of  the 
first-order  spectrum.  The  radiometric  measurements  of  the  dissoci- 
ation constants  of  the  indicators  were  all  made  between  X  =  0.5ju  and 
X  =  0.6/u,  and  in  this  region  the  first-order  spectrum  is  absolutely  pure. 

The  drumhead  which  rotated  the  grating  through  definite  angles  is  a 
large  one,  its  diameter  being  4  inches.  It  is  attached  to  a  very  care- 
fully constructed  screw,  so  designed  that  each  complete  revolution 
changes  the  wave-length  of  light  falling  on  the  slit  by  approximately 
500  A.  u.  The  drum  contains  500  divisions,  and,  all  things  considered, 
it  is  quite  probable  that  the  wave-length  settings  are  accurate  to  within 
one  or  two  Angstrom  units. 

The  spectroscope  is  of  the  Littrow  mounting,  the  same  lens  serving 
both  as  telescope  and  collimator.  It  is  provided  with  two  4-inch  lenses 
made  by  Brashear.  One  lens  has  a  focal  length  of  72  inches  and  is 
intended  for  photographic  work.  The  other,  and  the  one  which  was 
used  in  this  work,  has  a  focal  length  of  30  inches.  With  this  lens,  the 
visible  first-order  spectrum  at  slit  s2  has  a  length  of  over  6  inches. 
Provision  has  also  been  made  for  the  mounting  of  a  large  glass  prism 
on  the  grating  table,  and  it  is  purposed  to  use  this  prism  when  an 
intense  first-order  spectrum  is  desired,  rather  than  a  spectrum  which 
is  widely  dispersed.  Had  we  not  been  successful  in  constructing  a 
very  sensitive  radiomicrometer,  it  would  have  been  necessary  to  make 
frequent  use  of  the  glass  prism. 

The  spectroscope  was  placed  in  a  brass  box,  blackened  on  the  inside, 
which  prevented  stray  sources  of  light  from  reaching  the  grating.  The 
slits  were  of  the  bilateral  type. 


50  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

The  calibration  curve  for  the  grating  and  drumhead  was  calculated 
from  the  known  grating  space,  by  means  of  the  equation 

3000r+6fl 


oo  fiAfi  7  0^ 
X  =  00,800.7  sin 

where  T  is  the  number  of  complete  turns  of  the  screw  which  rotates  the 
grating  and  R  is  the  reading  on  the  drum.  The  calibration  was  also 
effected  by  observing  the  positions  of  various  mercury,  sodium,  and 
lithium  lines.  It  was  decided  that  the  calculated  values  are  more 
accurate  than  the  observed;  and,  accordingly,  the  dispersion  curve 
based  on  the  calculated  values  was  used. 

SOURCE  OF  LIGHT. 

A  Nernst  glower  served  as  the  source  of  light,  the  electrical  energy 
being  supplied  by  a  series  of  storage  batteries.  By  means  of  a  rheostat 
the  light  intensity  could  be  kept  quite  uniform.  The  glower  was  pro- 
tected from  air  drafts  by  means  of  a  box  of  asbestos  wood.  Provision 
was  made  for  the  mounting  of  a  nitrogen  lamp1  when  it  was  desired 
to  secure  large  deflections  of  the  radiomicrometer  for  wave-lengths  of 
light  shorter  than  X  =  0.5^- 

THE  CELLS. 

A  very  important  part  of  the  equipment  is  the  cells  which  contain 
the  solutions  to  be  investigated.  There  are  two  cells  made  as  nearly 
alike  as  possible.  The  cell  consists  of  two  brass  cylinders  which  closely 
telescope  into  one  another.  One  end  of  each  cylinder  is  closed  by  a 
glass  plate  held  in  position  by  Wood's  fusible  metal.  The  four  glass 
ends  used  in  the  cells  are  all  of  the  same  thickness,  i.  e.,  2  mm.  Their 
surfaces  are  plane,  and  are  parallel  to  within  5  wave-lengths  of  light. 
It  is  difficult  to  set  these  plates  in  the  brass  cylinder  with  Wood's 
fusible  metal,  without  warping  them  and  destroying  their  plane- 
parallelism.  The  character  of  the  interference  fringes  which  the  plates 
gave  with  the  mercury  arc  determined  when  they  were  correctly 
adjusted.  A  fine  thread  was  "chased"  on  the  outer  cylinder.  This 
thread  carried  a  nut  which,  when  turned,  raised  or  lowered  the  inner 
cylinder  a  definite  amount.  Each  complete  revolution  of  this  nut 
changed  the  distance  between  the  plates  by  1  mm.  The  nut  contained 
100  divisions,  and  by  means  of  this  arrangement  we  could  readily 
adjust  the  depth  of  the  solutions  to  within  less  than  0.01  mm.  Each 
cell  was  then  filled  with  a  layer  of  water  5  mm.  in  depth,  and  the 
deflections  given  by  the  radiomicrometer  were  noted  for  a  light  source 
of  uniform  intensity  throughout  the  whole  region  of  the  spectrum 
under  investigation,  namely,  from  X  =  0.4/x  to  X  =  2.0/1.  The  deflections 

'Dr.  W.  R.  Whitney,  of  the  General  Electric  Company,  very  kindly  supplied  us  with  two  lamps 
of  special  design  to  be  used  for  this  purpose. 


Radiometric  Measurements  of  Constants  of  Indicators. 


51 


should  agree  with  each  other  very  closely  for  all  wave-lengths  of  light 
if  the  cells  are  optically  identical.  Having  made  sure  that  the  cells 
are  optically  identical,  they  were  heavily  plated  with  gold  to  remove 
any  possibility  of  the  solutions  attacking  the  metals  of  the  cell.  It  is 
very  important  to  keep  the  solutions  perfectly  clear  when  measurements 
are  being  made,  and  the  glass  ends  must  be  maintained  absolutely  clean. 
Table  29  shows  the  optical  identity  of  the  two  cells.  Under  cell  A 
and  cell  B  are  given  the  actual  radiomicrometer  deflections  for  various 
wave-lengths  of  light. 

TABLE  29. — Comparison  of  the  two  cells. 


X=A.U. 

Cell  A. 

CellB. 

X  =  A.  U. 

Cell  A. 

CellB. 

4546 

4.0 

4 

8029 

125.0 

125.0 

4797 

8.2 

8.2 

8275 

127.0 

127.5 

5047 

14.2 

14.5 

8520 

129.0 

129.0 

5298 

22.0 

22.5 

9009 

129.5 

130.0 

5548 

31.5 

31.5 

9497 

115.5 

115.0 

5797 

41.5 

41.7 

9982 

105.5 

106.0 

6046 

50.5 

50.5 

10465 

113.5 

114.0 

6295 

60.7 

60.7 

10946 

108.5 

108.0 

6543 

70.5 

71.2 

11424 

69.0 

69.0 

6792 

86.7 

86.5 

11900 

51.5 

52.0 

7041 

100.5 

100.0 

12373 

56.0 

56.0 

7290 

107.7 

107.7 

12843 

53.5 

54.0 

7536 

114.0 

114.5 

13314 

34.0 

34.0 

7784 

120.5 

120.5 

13775 

11.0 

10.7 

Having  discussed  the  principal  parts  of  the  apparatus,  viz,  the  radio- 
micrometer,  spectroscope,  source  of  light,  and  the  cells,  and  having 
shown  how  they  meet  the  requirements  demanded  by  work  of  this 
character,  it  is  desirable  to  consider  next  the  general  arrangement  of 
the  various  parts. 

ARRANGEMENT  OF  APPARATUS. 

The  Nernst  glower,  lenses,  prisms,  and  carriage  for  the  cells  were 
mounted  on  an  upright  steel  standard  placed  by  the  side  and  near 
the  end  of  the  spectroscope  next  to  the  radiomicrometer.  Parallel 
light  from  the  Nernst  glower  was  made  to  pass,  by  means  of  a  carefully 
adjusted  sliding  carriage,  first  through  one  cell  and  then  through  the 
other.  The  two  cells  could  thus  be  made  to  occupy  the  same  position 
in  the  path  of  the  light.  The  light  transmitted  by  the  solutions  con- 
tained in  the  cells  was  then  deflected  by  means  of  a  right-angle  prism 
and  focused  on  the  slit  of  the  spectroscope  si.  Inside  of  the  box 
inclosing  the  spectroscope,  and  directly  in  the  rear  of  slit  «i,  was  placed 
another  right-angle  prism  which  reflected  the  light  through  a  4-inch 
lens  to  the  grating.  The  dispersed  light  was  reflected  back  through 
this  same  lens  and  focused  on  the  slit  of  the  spectroscope  s2.  The  light 
that  emerged  through  slit  s2  was  again  brought  to  a  focus  on  the  June- 


52  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

tion  of  the  radiomicrometer.  By  turning  the  drumhead  of  the  spectro- 
scope to  the  proper  points,  we  could  then  determine  quantitatively  the 
light  transmitted  by  any  solution  for  all  wave-lengths  of  light  between 


THE  DIFFERENTIAL  METHOD. 

It  was  desired  to  measure  the  absolute  percentage  transmission  of 
20  mm.  of  solution.  This  was  done  by  a  differential  method  which 
eliminated  corrections  for  reflection  from  the  glass  ends,  and  differences 
in  the  refractive  index  of  the  glass  and  the  solutions.  The  method  of 
procedure  is  as  follows:  Cell  A  is  filled  with  21  mm.  of  solution;  cell 
B  with  1  mm.  of  solution.  Light  of  unvarying  intensity  I0  was  then 
passed  through  cell  A,  and  the  intensity  of  the  transmitted  light  I, 
was  measured  by  means  of  the  radiomicrometer.  As  soon  as  possible, 
cell  B  was  made  to  occupy  the  same  position  formerly  occupied  by 
cell  A,  and  the  intensity  of  its  transmitted  light  I,,  was  determined. 
The  deflection  produced  when  cell  A  was  in  the  path  of  light,  divided  by 
that  given  when  cell  B  occupied  the  same  position,  determines  the 
absolute  percentage  transmission  for  20  mm.  of  solution;  or,  in  other 
words,  Ij/Iu  is  the  value  desired.  The  justification  for  this  procedure 
is  seen  from  what  follows. 

If  the  depth  of  solution  under  investigation  is  I,  the  intensity  of  the 
incident  light  I0,  and  that  of  the  transmitted  light  I,  we  have  the  follow- 
ing relations  for  depths  of  solution  I'  and  I"  : 

I,  =  Io  e-"  (a) 

!„  =  Io  e-""  (b) 

Dividing  a  by  b  we  have 

!,/!„  =  ^'"-v  (c) 

But  the  actual  percentage  transmission  of  the  same  solution  of  depth 
V  —  I"  =  I  is  given  by 

I/Io  =  e-k(l'~n  =  e-*(>"-l\  or  !,/!„  =  I/I0  (d) 

THEORETICAL  DISCUSSION. 

The  Ostwald1  theory  of  indicators,  explaining  first  the  cause  of  color, 
and  second  the  difference  in  sensitiveness  of  various  indicators  towards 
different  acids  and  bases,  has  been  found  to  be  inadequate.  Con- 
cerning the  latter,  which  is  by  far  the  most  important  side  of  the 
indicator  question,  the  Ostwald  interpretation  is  substantially  correct. 
But  in  consideration  of  well-known  relations  between  the  color  and 
structure  of  organic  compounds,  and  of  the  researches  of  Bernthsen,2 
Nietzi  and  Burckhart,3  Hantzsch,4  and  others,  it  has  been  found  neces- 

'Lehrbuch  der  all.  Chem.,  1,  799  (1891). 

'Chem.  Zeit,,  1956  (1892)  ;  also  Friedlander:  Ber.  d.  deutsch.  chem.  Gesell.,  26,  172,  2258  (1893). 

'Ibid.,  30,  175  (1897).  'Ibid.,  32,  583,  3085  (1899). 


Radiometric  Measurements  of  Constants  of  Indicators.  53 

sary  to  modify  the  Ostwald  view  as  to  the  cause  of  color.  The  facts 
and  relations  brought  out  by  these  investigations  have  been  correlated 
and  interpreted  by  Stieglitz1  in  the  so-called  chromophoric  theory  of 
color.  A.  A.  Noyes2,  in  a  quantitative  application  of  the  theory  of 
indicators  to  volumetric  analysis,  has  also  fully  explained  the  signifi- 
cance of  the  chromophoric  theory.  According  to  this  theory  it  is 
necessary  to  consider  an  indicator  solution  as  containing  a  mixture  of 
two  tautomeric  substances  of  different  structural  types.  The  ioni- 
zation  constants  of  the  two  forms,  and  the  equilibrium  relations  between 
them,  are  such  that  when  the  indicator  exists  as  a  slightly  ionized  acid 
or  base,  one  form  is  present  in  greatly  predominating  quantity.  The 
other  form  largely  predominates  when  the  indicator  exists  as  a  highly 
ionized  salt. 

Considering  phenolphthalein,  the  ionization  constant  K,,  according  to 
the  Ostwald  conception,  is  expressed  by  the  simple  equilibrium  equation 

HXP  =  K,XPH  (1) 

The  chromophoric  theory,  as  Stieglitz3  has  shown,  requires  two  such 
equations,  (a)  and  (6) : 

LHXfc  =  QH  (a) 

QXH  =  K'XQH  (6) 

where  k  is  the  stability  constant  expressing  the  equilibrium  relation 
between  the  two  tautomeric  acids.  The  acid  represented  by  LH  is 
assumed  to  be  a  pseudo-  or  an  extremely  weak  acid;  and  that  by  QH 
is  the  true  acid.  Its  ionization  constant  K'  is  of  such  a  magnitude  that 
the  quinoid  salt  is  formed  in  greatly  predominating  quantity  in  the 
presence  of  alkalies,  the  stability  constant  k  acting  so  as  to  maintain 
the  equilibrium  relation  between  the  two  tautomeric  acids.  The 
ionization  constant  for  phenolphthalein  is  the  product  of  the  stability 
constant  k,  and  the  ionization  constant  K'  of  the  acid  QH ;  or,  combining 
a  and  6  and  incorporating  k  and  K'  into  K,  we  have 

Q  X  H  =  K,X  LH  (2) 

The  above  equations  illustrate  the  fundamental  differences  between 
the  two  color  theories;  and  accepting  the  Stieglitz  interpretation  of  K, 
according  to  equation  2,  the  theory  underlying  the  calculation  of  the 
dissociation  constant  of  methyl  orange  from  radiometric  measurements 
will  be  discussed. 

Methyl  orange  is  in  reality  a  weak  base.  Noyes4  has  deduced  a 
general  expression  for  the  equilibrium  relations  of  any  pair  of  tauto- 
meric bases  and  their  ions.  This  deduction  involves  three  fundamental 

'Journ.  Amer.  Chem.  Soc.,  25,  1112  (1903).  *Itnd.,  32,  858  (1910). 

'Ibid.,  32,  815  (1910).  4/Wd.,  32,  818  (1910). 


54  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

equations.  Expressing  the  equilibrium  relation  k,  between  the  quinoid 
and  azo-base,  it  being  understood  that  the  symbols  represent  gram- 
molecular  or  gram-ionic  concentrations,  we  have 

QOH 
AzOH  ~ 

The  quinoid  base  and  the  azo-base  are  also  in  equilibrium  with  their 
ions,  according  to  4  and  5: 

QXOH 
QOH 

As  X  OH 
AzOH          KAZOH 

Multiplying  3  by  4  and  adding  5  to  the  product,  we  have 
Q  X  0"H  +  Az  X  OH 

AzOH  =  **Q03H     I      -tt-AzOH 

and  substituting  in  the  denominator  for  AzOH  its  value  —  - 
it  follows  that 

Q  X  OH  +  Az  X  OH       k  X  KQOH  +  KAzOH 
AzOH  +  QOH  1  +  k 

Letting  the  above  equal  K,-,  we  have 


__ 
AzOH  +  QOH  ~ 

Noyes  has  called  attention  to  the  fact  that  for  a  satisfactory  two-color 
indicator  such  as  methyl  orange,  the  sum  of  the  two  tautomeric  bases 

(AzOH  +  QOH)  must  be  substantially  equal  to  AzOH;  and  that  the 
+        +  + 

sum  of  the  two  ions  (Q  +  Az)  must  be  substantially  identical  with  Q. 

It  therefore  follows  that 

OH  X  Q 
AzOH 

where  K;  expresses  the  equilibrium  relations  of  the  two  tautomeric  bases 

and  their  ions,  and  is  substantially  the  equation  derived  by  Stieglitz.1 

+ 

Combining  equation  9  with  that  of  the  ion  product  of  water,  H  X  OH 
=  KK,  we  get 

H  X  AzOH       K. 

Q+  ~  ~K    :=  •l:N-(hydro1ysis) 

which  is  in  reality  the  familiar  equation  of  Walker2  (  -         ,        -  =  k  J  . 

'Journ.  Amer.  Chem.  Soc.,  25,  1112  (1903).  2Zeit.  phys.  Chem.,  4,  324  (1889). 


Radiometric  Measurements  of  Constants  of  Indicators.  55 

Having  determined  the  hydrolysis  constant  according  to  equation  10, 
the  ionization  constant  K;  of  methyl  orange  as  a  base  can  readily  be 

obtained.     The  procedure,  then,  is  to  determine  by  radiometric  meas- 

+  + 
urements,  the  concentrations  of  Q,  H,  and  AzOH  in  solutions  of  methyl 

orange  containing  varying  amounts  of  these  constituents.     A  method 

+ 

was  developed  whereby  the  concentration  of  the  quinoid  ions  Q  in 
equation  10  could  be  determined  from  the  light  transmitted  by  the 
indicator  solutions.  The  percentage  transmissions  for  these  solutions 
were  given  by  the  radiomicrometer  deflections.  The  method  will  be 
made  clear  by  the  following  theoretical  considerations  applied  to 
methyl  orange. 

If  we  consider  light  to  pass  through  an  absorbing  solution  of  depth  /, 
the  solvent  itself  having  no  absorption,  the  rate  of  change  of  intensity 
dl  is  given  by 

dl  =  -kltU  (11) 

The  constant  k  depends  only  on  the  wave-length  of  light  and  the 
nature  of  the  absorbing  medium.  If  I0  denotes  the  intensity  of  the 
incident  light,  then,  when  I  =  0,  I0  =  I  =  constant.  Integrating,  the 
intensity  of  the  transmitted  light  I,  given  by  an  absorbing  solution  of 
depth  I,  and  concentration  c,  is 

I  =  I0e-Wc  (12) 

If  the  solution  has  a  second  absorbing  component,  the  light  trans- 
mitted by  it  will  be 

It  =  I0  e-™  (13) 

Since  the  total  transmission  is  the  product  of  the  separate  trans- 
missions, we  have  for  the  actual  percentage  transmission  of  a  solution 
containing  two  absorbing  components  such  as  methyl  orange 

I/Io  =  e-klc  ~  k'l'c';  or  In  (I/I0)  =  -  kk  -  k'l'c'  (14) 

Since  the  depth  of  solution  was  maintained  constant  (20  mm.),  the 
above  equation  becomes 

In  (I/Io)  =  -  Kc  -  K'c'  (15) 

Applying  this  equation  to  methyl  orange,  let  c  represent  the  con- 
centration of  the  quinoid  salt  or  Q  in  equation  10,  and  c'  that  of  the 
azo-base.  If  T  is  the  total  quantity  of  methyl  orange  in  solution,  then 
c'  =  (T  —  c) ;  or,  since  a  dibasic  acid  was  used,  viz,  sulphuric  acid, 
c'  =  (T  — 2c).  When  a  pure  solution  of  methyl  orange  is  slightly 
acidified  with  sulphuric  acid,  and  not  all  of  the  azo-base  converted  into 
the  quinoid  salt,  both  c  and  c'  are  present.  The  light  transmitted  by 
such  a  solution  will  be: 

In  (I/Io)  =  -  Kc  -  K'(T  -  2c)  (16) 


56  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

In  a  pure  aqueous  solution  of  methyl  orange,  or  one  containing  an 
excess  of  alkali,  c  =0,  therefore,  equation  15  reduces  to: 

ln(I/I0)'=  -KY=  -K'T  (17) 

If  to  a  solution  of  methyl  orange  sufficient  sulphuric  acid  is  added  to 
convert  all  of  the  azo-base  into  the  quinoid  salt  and  completely  suppress 
hydrolysis,  c'  =  0,  and  equation  15  becomes 

In  (I/I0)"  =  --  Kc  =  --  KT/2  (18) 

From  equations  16,  17  and  18,  K  and  K'  can  be  eliminated,  and 
solving  for  c  we  obtain: 

T[m(I/I0)    -In  (I/Io)'] 
-  2[ln  (I/Io)"  -  In  (I/Io)'] 

In  the  above  equation  (I/Io)'  is  the  percentage  transmission  for  the 
solution  of  pure  methyl  orange  for  some  given  wave-length  of  light; 
(I/I0)",  the  percentage  transmission  of  the  solution  containing  an  excess 
of  acid  for  the  same  wave-length,  and  (I/Io)  the  percentage  transmission 
for  the  same  wave-length  of  the  solution  whose  quinoid  salt  concen- 
tration c  is  to  be  determined. 

Returning  now  to  the  fundamental  hydrolysis  equation  for  methyl 
orange  previously  derived : 

H  X  AzOH       K,, 
Q+  =  K, 

we  can  readily  insert  the  proper  values,  knowing  the  total  concentration 
of  methyl  orange  T,  and  having  determined  by  radiometric  means  the 

quinoid  salt  concentration  c.  Q  is  equal  to  2c.  AzOH  =  c'  —  T  —  2c. 
+ 

H  is  given  by  2(T'--  c),  where  T'  represents  the  total  quantity  of  acid 
added.  It  is  assumed  that  the  dissociation  of  these  extremely  dilute 
solutions  is  practically  complete.  The  sulphonic  acid  group  can  have 
but  little  effect  on  the  hydrogen  ion  concentration,  since  benzenesul- 
phonic  acid1  is  as  strong  as  sulphuric  acid,  being  dissociated  at  25°  to  the 
extent  of  90  per  cent  for  a  dilution  v  =32.  As  Stieglitz2  has  pointed 
out,  the  whole  behavior  of  methyl  orange  is  that  of  a  very  weak  base, 
and  the  elimination  of  the  sodium  sulphonate  group  from  it  leaves 
dimethylaniline  azobenzene,  which  shows  all  the  characteristics  of 
methyl  orange  as  an  indicator. 

It  will  be  noticed  from  equation  10  that  it  is  necessary  to  know  the 
ionization  constant  for  water  before  the  constant  for  the  indicator  can 
be  calculated  from  its  hydrolysis  constant.  The  generally  accepted 
value  of  Kw  at  25°  is  1.2X10"14.  It  was  desired  to  know  the  value  of 
the  constant  at  20°,  since,  unless  otherwise  stated,  it  was  at  this  tem- 

'Carnegie  Inst.  Wash.  Pub.  No.  170,  128  (1912). 
"Journ.  Amer.  Chem.  Soc.,  25,  1117  (1903). 


Radiomelric  Measurements  of  Constants  of  Indicators.  57 

perature  that  all  measurements  were  made.  Owing  to  the  large  value 
of  the  heat  of  ionization  of  water,  the  value  of  its  ionization  constant 
is  much  less  at  20°  than  at  25°.  Just  what  this  change  is  may  be  cal- 
culated from  the  well-known  Van't  Hoff  formula 

Ki       g  (T.  -  TJ 
a  K2       R(T2  X  TO 

where  KI  and  K2  represent  the  ionization  constants  for  water  at  tem- 
pera turesTi  and  T2.  RI  is  the  gas  constant,  which  equals  1.986  calories 
per  degree,  and  q  is  the  heat  of  ionization  of  water  or  13,700  calories. 
Inserting  these  values  in  equation  20,  KW  =  0.81X10~14  at  20°.  This 
value  was  used  for  the  calculation  of  the  ionization  constants  of  both 
methyl  orange  and  phenolphthalein. 

EXPERIMENTAL  WORK  ON  METHYL  ORANGE. 

Before  discussing  the  tables  showing  the  results  of  the  calculations 
based  on  the  above  deductions,  it  will  perhaps  be  of  interest  to  consider 
a  few  of  the  more  important  results  of  the  preliminary  work  on  methyl 
orange.  Two  mother  solutions  of  known  concentrations  were  pre- 
pared, the  one  being  methyl  orange  and  the  other  sulphuric  acid. 
All  solutions  were  made  up  at  20°,  and  carefully  purified  substances 
dissolved  in  conductivity  water  were  employed  in  all  cases.  A  number 
of  test  solutions  were  prepared  from  the  mother  solutions,  all  of  which 
contain  equal  amounts  of  methyl  orange  but  different  amounts  of  sul- 
phuric acid.  The  volume  of  each  solution  was  100  c.c.  The  solutions 
thus  presented  a  series  of  color  shades,  ranging  from  yellow  to  deep  red. 
The  percentage  transmissions  I/Io,  were  taken  with  a  20  mm.  depth  of 
each  solution  for  the  same  4  or  5  wave-lengths  of  light.  The  region 
of  the  spectrum  to  be  studied  is  given  by  the  ascending  arm  of  the  trans- 
mission curve  for  methyl  orange.  This  region  for  the  above-named 
indicator  is  between  X  =0.56^  and  X  =  0.59/i.  The  radiomicrometer 
deflections  in  this  region  are  necessarily  small,  and  certain  variations 
which  appear  in  the  data  can  be  explained  in  a  large  measure  as  due  to 
the  vibrations  of  the  building,  which  often  prevented  accurate  readings. 

Special  attention  is  called  to  a  solution  of  methyl  orange  containing 
an  excess  of  alkali,  and  another  solution  containing  an  excess  of  acid. 
Equation  19  shows  that  the  calculation  of  c  depends  not  only  on  the 
percentage  transmission  of  the  solution  in  question,  but  also  on  the 
percentage  transmission  for  a  solution  in  which  all  of  the  methyl  orange 
exists  as  the  azo-base,  and  one  in  which  all  of  the  indicator  has  been 
converted  into  the  quinoid  salt.  It  is  thus  necessary  to  know  if  any 
alkali  must  be  added  to  convert  all  of  the  methyl  orange  into  the  azo- 
base  and  to  prevent  hydrolysis;  and  also  how  much  acid  is  required  to 
form  the  quinoid  salt  and  to  suppress  hydrolysis  completely.  Further- 
more, the  stability  of  these  solutions  must  be  considered. 


58 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


Table  30  shows  that  in  a  pure  aqueous  solution  of  methyl  orange 
there  is  no  appreciable  hydrolysis,  and  that  practically  all  of  the  indi- 
cator exists  as  the  azo-base.  It  is,  therefore,  not  necessary  to  add  alkali 
in  determining  the  value  of  (I/Io)'  in  equation  19.  The  volume  of  all 
solutions  used  in  table  30  was  100  c.c.  and  each  solution  contained  the 
same  amount  of  methyl  orange. 

TABLE  30. 
[I/Io  for  depth  of  solution  =  20  mm.] 


X  =A.  U. 

Methyl  orange 
in  pure  water. 

Methyl  orange  plus 
1  c.c.  N  sodium 
hydroxide. 

Methyl  orange  plus 
2.5  c.c.  N  sodium 
hydroxide. 

5648 

70.3 

71.2 

70.5 

5698 

79.3 

78.8 

79.0 

5748 

84.6 

86.0 

84.8 

5797 

90.0 

90.0 

88.5 

5847 

91.4 

92.2 

90.0 

Table  31  gives  results  for  three  solutions  of  methyl  orange,  each  con- 
taining an  excess  of  acid  and  the  same  amount  of  indicator  diluted  to 
100  c.c.  It  shows  that  0.2  c.c.  of  the  acid  is  sufficient  to  convert  all  of 
the  azo-base  into  the  quinoid  salt  and  completely  to  suppress  hydrolysis. 

TABLE  31. 
[I/Io  for  depth  of  solution  =  20  mm.] 


A=A.  U. 

Methyl  orange  plus 
0.2  c.c.  concentrated 
sulphuric  acid. 

Methyl  orange  plus 
0.5  c.c.  concentrated 
sulphuric  acid. 

Methyl  orange  plus 
1  .0  c.c.  concentrated 
sulphuric  acid. 

5698 

15.2 

15.2 

16.6 

5723 

20.2 

20.1 

20.1 

5748 

30.7 

31.2 

31.9 

5773 

39.5 

39.2 

38.8 

5797 

48.7 

49.2 

48.1 

5823 

55.0 

53.8 

55.3 

5847 

61.4 

62.6 

63.2 

Stable  solutions  of  methyl  orange  are  given  only  when  the  concen- 
tration is  less  than  2  X  10~4  gram-molecules  per  liter.  Considerable 
difficulty  was  encountered  during  the  preliminary  work,  owing  to  the 
fact  that  the  indicator  solutions  were  more  concentrated  than  the  above- 
named  limit.  Such  solutions  of  methyl  orange,  especially  those  con- 
taining a  large  excess  of  acid  and  also  those  with  enough  acid  to  give 
an  intense  red  color,  gradually  become  more  transparent  on  standing. 
The  action  of  light  very  greatly  accelerates  this  bleaching  process. 
The  solutions  most  susceptible  to  the  action  of  light  and  of  time  are 
those  containing  the  greatest  amounts  of  acid.  It  was  found  that  these 
solutions  could  be  kept  in  the  dark  for  4  or  5  hours  unchanged,  and  that 
bleaching  again  took  place  as  soon  as  they  were  exposed  to  the  light. 
Experiments  were  made  which  showed  that  light  of  the  shorter  wave- 
lengths was  more  actinic  in  increasing  the  transparency  of  the  solutions 


Radiometric  Measurements  of  Constants  of  Indicators.  59 

than  light  of  longer  wave-lengths.  The  addition  of  10  per  cent  of  ethyl 
alcohol  rendered  these  solutions  fairly  stable  to  the  action  of  light 
and  time.  After  standing  for  a  day  or  two,  solutions  containing  an 
excess  of  acid  often  showed  the  presence  of  fine  crystals.  Large 
amounts  of  these  crystals  could  be  obtained  by  strongly  acidifying 
a  saturated  solution  of  methyl  orange.  It  is  thought  that  the  forma- 
tion of  these  crystals,  very  probably  helianthine  itself,  causes  the 
instability  of  the  solutions  in  question.  The  obvious  procedure  to 
follow  was  to  use  more  dilute  solutions,  which  were  found  to  be  quite 
stable,  and  even  the  solution  containing  the  excess  of  acid  would  remain 
practically  unchanged  when  exposed  to  the  action  of  light  for  a  short 
time.  However,  when  the  concentration  of  methyl  orange  approxi- 
mated 2X10~4  gram-molecules  per  liter,  as  is  the  case  for  the  solutions 
given  in  table  32,  the  solution  containing  an  excess  of  acid  was  pre- 
pared last  and  its  percentage  transmission  measured  immediately. 
The  remaining  solutions  were  kept  in  the  dark  until  used.  It  is  certain 
that  when  such  precautions  are  taken,  no  appreciable  change  in  trans- 
parency could  occur  before  the  radiometric  measurements  were  com- 
pleted. Table  32,  the  concentration  of  the  methyl  orange  being  1.98  X 
10~4  gram-molecules  per  liter,  indicates  that  even  the  solutions  con- 
taining an  excess  of  acid  can  be  safely  used  in  this  way. 

The  following  shows  how  accurately  small  concentrations  of  hydro- 
gen and  hydroxyl  ions  can  be  estimated  by  radiometric  measurements. 
If  the  percentage  transmission  for  the  various  solutions  be  plotted  as 
curves,  the  abscissa}  being  wave-lengths  and  the  ordinates  percentage 
transmissions,  it  will  be  noted  that  with  increasing  amounts  of  acid  the 
transmission  curve  for  pure  methyl  orange  is  widely  displaced  towards 
the  red  end  of  the  spectrum.  The  solution  curve  corresponding  to 
the  excess  of  acid  is  displaced  about  400  A.  u.  from  the  curve  corre- 
sponding to  a  pure  aqueous  solution  of  methyl  orange.  From  the 
displacements  produced  by  solutions  containing  known  amounts  of 
acid,  and  the  displacement  given  by  a  solution  containing  an  unknown 
amount,  the  concentration  of  the  unknown  amount  of  acid  can  be 
quite  accurately  and  quickly  determined.  In  one  case  a  solution  was 
prepared  containing  an  amount  of  sulphuric  acid  unknown  to  us.  The 
concentration  of  the  acid  was  determined  by  the  above  method,  and 
the  value  found  was  0.00000216  gram  per  cubic  centimeter.  The 
amount  actually  present  was  0.00000210  gram  per  cubic  centimeter. 
Several  similar  attempts  were  made  and  they  were  quite  as  successful. 

RESULTS  WITH  METHYL  ORANGE. 

The  hydrolysis  and  ionization  constants  for  methyl  orange,  recorded 
in  Tables  33  and  35,  are  to  be  regarded  as  part  of  the  preliminary  work. 
Table  32  contains  the  percentage  transmissions  for  nine  solutions  of 
methyl  orange  prepared  in  accordance  with  the  scheme  given  below. 
The  concentration  of  the  mother  solutions  of  methyl  orange  was  3.973 


60 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


XlO  4;  that  of  the  mother  solution  of  sulphuric  acid  3.0885X10" 
gram-molecules  per  liter. 

SOLUTIONS. 

I.  50  c.c.  methyl  orange,  diluted  to  100  c.c. 
II.  50  c.c.  methyl  orange,  5  c.c.  sulphuric  acid,  diluted  to  100  c.c. 

III.  50  c.c.  methyl  orange,  10  c.c.  sulphuric  acid,  diluted  to  100  c.c. 

IV.  50  c.c.  methyl  orange,  15  c.c.  sulphuric  acid,  diluted  to  100  c.c. 
V.  50  c.c.  methyl  orange,  20  c.c.  sulphuric  acid,  diluted  to  100  c.c. 

VI.  50  c.c.  methyl  orange,  30.9  c.c.  sulphuric  acid,  diluted  to  100  c.c. 

VII.  50  c.c.  methyl  orange,  35  c.c.  sulphuric  acid,  diluted  to  100  c.c. 

VIII.  50  c.c.  methyl  orange,  1.0  c.c.  cone,  sulphuric  acid,  diluted  to  100  c.c. 

IX.  50  c.c.  methyl  orange,  2.0  c.c.  cone,  sulphuric  acid,  diluted  to  100  c.c. 

TABLE  32. 
[I/Io  for  depth  of  solution  =20  mm.] 


X=A.  U. 

I. 

II. 

III. 

IV. 

V. 

VI. 

VII. 

VIII. 

IX. 

Average  of 
VIII  and  IX. 

5748 

84.6 

71.9 

65.1 

57.7 

49.1 

46.7 

9.5 

8.4 

8.95 

5798 

90.0 

85.6 

79.3 

73.3 

70.0 

62.2 

61.0 

22.1 

22.1 

22.1 

5847 

91.4 

84.6 

81.4 

77.9 

70.3 



35.7 

35.7 

35.7 

Attention  is  called  to  the  blank  spaces  appearing  in  some  of  the  tables. 
For  some  solutions  and  certain  wave-lengths  of  light  no  transmission 
values  were  obtained.  It  was  known  at  the  time  the  measurements 
were  made  that  certain  transmission  values  were  erroneous,  due  either 
to  variations  in  the  current  intensity  or  to  vibrational  disturbances. 
The  action  of  light,  as  has  been  previously  explained,  renders  these 
solutions  of  methyl  orange  more  transparent,  and  for  this  reason  it 
was  necessary  to  complete  the  reading  as  soon  as  possible.  This 
excluded  the  possibility  of  remeasuring  the  percentage  transmissions 
made  when  conditions  were  unfavorable  for  accurate  readings. 

Table  33  gives  the  quinoid-salt  and  azo-base  concentrations  of  each 
of  the  above  solutions,  the  hydrolysis  constant  (K^/K,),  and  the  ioni- 
zation  constant  K;  of  methyl  orange  as  a  base. 

TABLE  33. 


Solution. 

X=A.U. 

Azo-base 
c'XIO*. 

Quinoid  salt. 
cXIO6. 

KWKiXlO'. 

K.'XIO11. 

II 

5798 

1.91 

0.353 

6.5 

1.3 

III 

5748 

.84 

0.716 

6.1 

1.3 

5797 

.80 

0.897 

4.4 

1.8 

5847 

.82 

0.813 

5.1 

1.6 

IV 

5748 

.75 

1.16 

6.2 

1.6 

5797 

.69 

1.46 

3.7 

2.2 

5847 

.74 

1.23 

4.8 

1.7 

V 

5748 

1.64 

1.69 

4.4 

1.9 

5797 

1.63 

1.78 

4.0 

2.0 

5847 

1.65 

1.68 

4.4 

1.8 

VI 

5748 

1.50 

2.41 

4.5 

1.8 

5797 

1.46 

2.61 

3.9 

2.0 

5847 

1.43 

2.76 

3.5 

2.3 

VII 

5847 

1.46 

2.62 

4.6 

1.8 

5797 

1.46 

2.65 

4.5 

1.8 

Av.,4.5X10-« 

Av.,  1.8X10-" 

Radiometric  Measurements  of  Constants  of  Indicators. 


61 


The  calculations  of  the  hydrolysis  and  ionization  constants  are  based 
on  logarithmic  functions;  and  where  the  transmission  of  the  solution  is 
nearly  complete,  a  slight  variation  from  the  actual  percentage  trans- 
mission will  cause  a  great  change  in  the  calculated  hydrolysis  and  ioni- 
zation constants.  If  we  consider  an  experimental  case  where  the 
amount  of  absorption  to  be  measured  is  slight,  the  above  can  be  made 
quite  clear.  Solution  II  in  table  32  for  wave-length  of  light  X  =  5798 
A.  u.  actually  gave  a  percentage  transmission  of  85.6.  Changing 
the  percentage  transmission  by  small  amounts,  and  calculating  the 
value  of  the  hydrolysis  and  ionization  constants  from  these  values,  we 
have 

I/Io.  K«,/Ki.  Ki. 

85.6  6.46X10-4  1 . 25  X  10-"  found  experimentally 

86.0  7.21X10-4  1. 12X10-" 

86.6  8.59X10-4  0. 94X10-" 

It  will  thus  be  seen  that  changing  the  percentage  transmission  by 
about  1  per  cent  will  cause  a  variation  in  the  ionization  constant  of 
about  25  per  cent.  When  conditions  are  not  satisfactory,  it  is  quite 
possible  that  the  error  in  some  of  the  percentage  transmissions  may  be 
greater  than  1  per  cent.  Considering  these  facts,  the  constants  show 
that  the  efforts  which  were  made  to  eliminate  errors  arising  from  varia- 
tions of  the  current  intensity  and  vibrational  disturbances  were  fairly 
successful. 

The  transmission  values  given  in  table  34  are  for  a  much  more  dilute 
solution  of  methyl  orange.  The  concentration  of  the  mother  solution 
of  methyl  orange  was  1.9865X10"4  gram-molecules  per  liter;  that 
of  the  mother  solution  of  sulphuric  acid  being  3.0885  X  10~4  gram- 
molecules  per  liter. 

SOLUTIONS. 

I.  50  c.c.  methyl  orange,  diluted  to  100  c.c. 
II.  50  c.c.  methyl  orange,  10  c.c.  sulphuric  acid,  diluted  to  100  c.c. 

III.  50  c.c.  methyl  orange,  15  c.c.  sulphuric  acid,  diluted  to  100  c.c. 

IV.  50  c.c.  methyl  orange,  20  c.c.  sulphuric  acid,  diluted  to  100  c.c. 

V.  50  c.c.  methyl  orange,  1  c.c.  cone,  sulphuric  acid,  diluted  to  100  c.c. 

VI.  50  c.c.  methyl  orange,  0.5  c.c.  cone,  sulphuric  acid,  diluted  to  100  c.c. 

VII.  50  c.c.  methyl  orange,  0.2  c.c.  cone,  sulphuric  acid,  diluted  to  100  c.c. 

VIII.  50  c.c.  methyl  orange,  0.5  c.c.  cone,  sulphuric  acid,  diluted  to  100  c.c. 

TABLE  34. 
[I/Io  for  depth  of  solution  =  20  mm.] 


X  -A.  U. 

I. 

II. 

III. 

IV. 

V. 

VI. 

VII. 

VIII. 

Average  V,  VI, 
VII  and  VIII. 

5698 

87.7 

75.3 

67.6 

66.3 

16.6 

15.2 

15.2 

17.1 

16.0 

5723 

88.9 

80.7 

74.2 

72.3 

20.1 

20.1 

20.2 

20.1 

5748 

90.3 

83.7 

79.8 

76.6 

31.9 

31.2 

30.7 

30.3 

31.3 

5773 

91.6 

81.7 

81.3 

38.8 

39.0 

39.5 

41.1 

39.9 

5797 

93.5 

88.9 

85.1 

84.4 

48.1 

49.2 

48.7 

47.5 

48.4 

5823 

94.5 

91.3 

87.0 

86.2 

55.3 

53.8 

55.0 

56.4 

55.1 

5847 

95.6 

91.7 

88.2 

63.2 

62.6 

61.4 

61.7 

62.2 

62 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


The  constants  given  in  table  35  are  calculated  from  the  values  in 
table  34. 

The  ionization  constants  for  methyl  orange,  recorded  in  tables  33  and 
35,  do  not  by  any  means  represent  the  most  accurate  values  that  can  be 
obtained  by  this  radiometric  method.  It  is  necessary  that  the  meas- 
urements be  made  in  a  region  of  the  spectrum  where  the  light  energy  is 
not  intense,  and,  consequently,  the  deflections  of  the  radiomicrometer 
are  small.  The  largest  deflection  that  could  be  obtained  for  wave- 
lengths of  light  X  =  0.58  fj,  was  60  mm.  It  is  for  this  reason  that  vibra- 
tions, and  variations  in  the  intensity  of  the  Nernst  glower,  seriously 
interfere  with  the  accuracy  of  the  percentage  transmissions.  The 
temperature  which  prevailed  at  the  time  these  measurements  were 
made  was  not  accurately  known.  However,  it  is  certain  that  it  was 
between  20°  and  23°.  The  constants  were  calculated  for  20°;  there- 
fore, there  is  good  reason  to  suppose  that  the  real  constants  are  some- 
what higher  (see  table  37). 

TABLE  35. 


Solution. 

X=A.  U. 

Azo-hase 
c'XIO5. 

Quinoid  salt. 
cXIO6. 

Ku./K,X104. 

KiXlO". 

IV 

5698 

8.30 

8.14 

5.5 

.5 

5723 

8.55 

6.88 

6.8 

.2 

5748 

8.41 

7.62 

6.0 

.4 

5773 

8.49 

7.18 

6.4 

.3 

5797 

8.42 

7.56 

6.0 

.4 

5823 

8.22 

8.56 

5.1 

.6 

5847 

8.06 

9.36 

4.5 

.8 

III 

5698 

8.42 

7.58 

4.3 

.9 

5723 

8.72 

6.03 

5.8 

.4 

5748 

8.78 

5.76 

6.2 

.3 

5773 

8.55 

6.88 

4.9 

.7 

5797 

8.54 

6.95 

4.8 

.7 

5823 

8.40 

7.63 

4.3 

.9 

II 

5698 

9.04 

4.47 

5.3 

.5 

5723 

9.28 

3.24 

7.9 

.0 

5748 

9.22 

3.53 

7.1 

.1 

5797 

9.20 

3.62 

6.9 

.2 

5823 

9.28 

3.22 

8.0 

.0 

5847 

8.96 

4.84 

4.8 

.7 

Av.,  5.8X10-4 

Av.,     .5X10-" 

The  transmission  values  given  in  table  36  were  obtained  when  the 
temperature  at  the  time  of  measurement  was  known  to  be  20°.  The 
measurements  were  made  at  night,  when  the  building  was  fairly  free 
from  vibrations.  Under  these  conditions,  the  scale  could  be  read 
accurately  to  0.25  mm.  Except  for  solutions  II  and  VI,  the  current 
intensity  remained  very  constant.  A  different  mother  solution  of 
methyl  orange  was  used,  its  concentration  being  3.054  X  10~4  gram- 
molecules  per  liter.  The  concentration  of  the  mother  solution  of  sul- 
phuric acid  was  3.089  X  10~4  gram-molecules  per  liter. 


Radio-metric  Measurements  of  Constants  of  Indicators. 


63 


I.  50  c.c.  methyl  orange, 
II.  60  c.c  methyl  orange, 

III.  50  c.c.  methyl  orange, 

IV.  50  c.c.  methyl  orange, 
V.  50  c.c.  methyl  orange, 

VI.  50  c.c.  methyl  orange, 

VII.  50  c.c.  methyl  orange, 

VIII.  50  c.c.  methyl  orange, 


SOLUTIONS. 

diluted  to  100  c.c. 

10  c.c.  sulphuric  acid,  diluted  to  100  c.e. 
15  c.c.  sulphuric  acid,  diluted  to  100  c.c. 
20  c.c.  sulphuric  acid,  diluted  to  100  c.c. 
25  c.c  sulphuric  acid,  diluted  to  100  c.c. 
40  c.c.  sulphuric  acid,  diluted  to  100  c.c. 
0.7  c.c.  cone,  sulphuric  acid,  diluted  to  100  c.c. 
0.8  c.c.  cone,  sulphuric  acid,  diluted  to  100  c.c. 


TABLE  36. 
[I/Io  for  depth  of  solution  =20  mm.     Temperature  =20°.] 


X=A.  U. 

I. 

II. 

III. 

IV. 

V. 

VI. 

VII. 

VIII. 

Average  VII 
and  VIII. 

5723 

88.2 

71.8 

63.7 

58.0 

51.6 

11  8 

11  6 

11  7 

5748 

90.0 

77.6 

68.6 

64.0 

59.8 

17.1 

17  6 

17  4 

5773 

91.6 

75.2 

71.0 

67.4 

24  8 

24  9 

24  9 

5797 

93  4 

78.7 

76.2 

72  6 

63  8 

32  5 

33  2 

32  4 

5848 

95.4 

89.6 

86.6 

84.5 

75  2 

49  0 

49  8 

49  4 

5898 

97  2 

91.4 

89.8 

88  7 

63  5 

63  5 

63  5 

The  values  calculated  from  table  36  are  recorded  in  table  37. 

TABLE  37. 


Solution. 

X=A.  U. 

Azo-base 
c'XIO4. 

Quinoid  salt 
cXIO6. 

Ktt/K,X104. 

KtXUf1. 

VI 

5797 

0.975 

2.76 

3.4 

2.4 

5848 

0.975 

2.76 

3.4 

2.4 

V 

5723 

1.193 

1.67 

4.3 

1.9 

5748 

1.147 

1.99 

3.5 

2.3 

5773 

1.167 

1.80 

3.8 

2.1 

5797 

1.203 

1.62 

4.5 

1.8 

5898 

1.197 

1.65 

4.4 

1.8 

IV 

5723 

1.211 

1.58 

3.5 

2.3 

5748 

1.211 

1.58 

3.5 

2.3 

5773 

1.227 

1.50 

3.8 

2.1 

5798 

1.235 

1.46 

4.0 

2.0 

5848 

1.315 

1.36 

4.7 

1.8 

5898 

1.245 

1.42 

4.2 

1.9 

III 

5723 

1.281 

1.23 

3.5 

2.3 

5748 

1.277 

1.25 

3.5 

2.3 

5773 

1.295 

1.16 

3.9 

2.1 

5798 

1.269 

1.24 

3.5 

2.3 

5848 

1.309 

1.09 

4.3 

1.9 

5898 

1.301 

1.13 

4.0 

2.0 

II 

5723 

1.372 

0.773 

4.1 

2.0 

5748 

1.390 

0.687 

4.9 

1.7 

5848 

1.383 

0.718 

4.6 

1.8 

Av.,  4.0X10-4 

Av.,  2.1X10-" 

The  greatest  variation  from  the  mean  is  less  than  20  per  cent.  It  has 
already  been  shown,  in  one  experimental  case,  that  an  error  of  1  per 
cent  in  the  percentage  transmission  would  cause  a  variation  of  25  per 
cent  in  the  ionization  constant.  The  determination  of  the  percentage 
transmissions  from  which  the  constants  recorded  in  table  37  were  cal- 


64  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

culated,  was  made  under  the  most  favorable  conditions.  These  con- 
stants, therefore,  more  likely  represent  the  true  constants  of  the 
indicator. 

PHENOLPHTHALEIN. 

Regarding  phenolphthalein  as  a  monobasic  acid,  we  have  from 
equation  2 

QxH 

~nr 

Dividing  this  by  the  ion  product  of  water,  it  follows  that 

LH  X  OH       Ktt 
~Q+~        =  K7 

Phenolphthalein  has  but  one  component  of  absorption,  and  the  appli- 
cation of  Beer's  law  is  much  simpler  than  for  methyl  orange  where 
there  are  two  components  of  absorption.  A  solution  of  phenolphthalein 
in  pure  water  is  perfectly  transparent.  The  addition  of  an  excess  of 
sodium  hydroxide  converts  all  the  colorless  lactoid  molecules  into  the 
red  quinoid  salt,  the  concentration  of  which  we  will  represent  by  c'. 
Let  c  be  the  concentration  of  the  quinoid  salt  in  the  phenolphthalein 
solutions  containing  ammonium  hydroxide  and  ammonium  chloride. 
Since  the  depth  of  solution  was  maintained  constant,  we  have  the  two 
fundamental  equations  given  below: 

In  (I/I0)  =  -  Kc  (22) 

ln(I/I0)'=-Kc'  (23) 

The  constant  K  is  the  same  in  both  cases.  Dividing  22  by  23  we 
have 


)  =  C/C'  or  c  =  c/  X  l  (24) 

ln(I/I0)'  In  (I/Io)' 

Where  (I/Io)'  is  the  percentage  transmission  of  a  solution  of  phenol- 
phthalein containing  an  excess  of  alkali,  for  some  given  wave-length 
of  light;  (I/Io)  that  for  the  solution  in  which  the  concentration  of  the 
quinoid  salt  c,  is  to  be  determined  for  the  same  wave-length  of  light. 
Knowing,  then,  the  amount  of  phenolphthalein  converted  into  the 
quinoid  salt,  LH  in  equation  21  is  given  by  T—  c,  where  T  represents 
the  total  amount  of  phenolphthalein.  The  hydroxyl  ion  concentration 
was  varied  by  the  addition  of  ammonium  chloride  to  ammonium 
hydroxide. 

The  value  of  OH  to  be  substituted  in  the  hydrolysis  equation  for 
phenolphthalein  was  obtained  from  the  following  equation  : 

NH4  X  O~H 


(NH4OH  +  NH3) 


Radiometric  Measurements  of  Constants  of  Indicators.  65 

If  ammonium  chloride  and  ammonium  hydroxide  are  present  in  the 
same  solution,  the  concentration  of  the  NKU  ions  is  furnished  almost 
entirely  by  the  ammonium  chloride.  Let  S  equal  the  concentration  of 
the  salt  and  B  the  concentration  of  the  base.  We  then  have  for  dilute 
solutions 

=  K,  (25) 


or 

OH  =  ^-|p  (26) 

The  ionization  constant  for  ammonium  hydroxide1  at  25°  is  given  as 
18X10'6.  All  of  the  work  on  phenolphthalein  was  done  at  20°,  and 
reducing  the  above  value  to  20°,  we  obtain  from  equation  20,  q  being 
equal  to  1,400  calories,  K6  =  17.4  X  10"6. 

A  pure  sample  of  phenolphthalein  was  recrystallized  from  absolute 
methyl  alcohol.  A  weighed  amount  of  purified  phenolphthalein  was 
dissolved  in  50  c.c.  of  absolute  ethyl  alcohol.  By  means  of  a  small  pipette, 
carefully  calibrated,  a  small  volume  of  this  alcoholic  solution,  say  1.25 
c.c.,  could  be  quite  accurately  measured.  This  volume  was  diluted 
with  conductivity  water  to  2,000  c.c.  Experiments  were  made  which 
showed  that  the  effect  of  this  small  trace  of  alcohol  was  negligible. 
(See  McCoy.)  The  ammonia  employed  for  the  solution  of  ammonium 
hydroxide  was  distilled  from  barium  hydroxide  to  eliminate  carbon 
dioxide.  All  solutions  were  prepared  with  conductivity  water  at  20°. 
Special  care  was  taken  to  prevent  carbon  dioxide  from  coming  in 
contact  with  any  of  the  solutions. 

It  is  well  known  that  an  excess  of  alkali  causes  a  rather  rapid  fading 
of  solutions  of  phenolphthalein.  Experiments  were  made  which  showed 
that  a  very  slight  error  would  be  introduced  if  the  solutions  of  phenol- 
phthalein containing  an  excess  of  alkali  were  used  directly  after  being 
prepared.  Filling  the  cells  with  the  solution,  and  taking  radiometric 
measurements  for  5  wave-lengths  of  light,  required  about  8  minutes, 
and  during  this  short  space  of  time  the  change  in  transparency  was 
found  to  be  very  slight. 

A  few  of  the  experiments  made  with  solutions  of  phenolphthalein 
containing  an  excess  of  alkali  are  recorded  in  table  38.  All  solutions 
were  diluted  to  100  c.c.  and  each  contained  equal  amounts  of  phenol- 
phthalein. A  normal  solution  of  sodium  hydroxide  was  used  for  these 
experiments.  The  percentage  transmissions  given  by  solutions  I  and 
II  are  for  0.8  c.c.  and  1  c.c.,  respectively,  of  the  alkali.  The  interval 
which  elapsed  between  the  time  the  solutions  were  prepared  and  the 
last  measurement  made  was  between  7  and  8  minutes.  Measurements 

'Carnegie  Inst.  Wash.  Pub.  No.  63,  298  (1907). 


66 


Conductitdties  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


were  again  made  with  solution  II,  15  minutes  later.  The  results  are 
recorded  under  III.  By  comparing  the  transmission  values,  the  effect 
of  the  bleaching  for  an  interval  of  15  minutes  can  be  noted.  Solution 
IV  contained  0.5  c.c.  alkali,  and  was  kept  in  the  dark  nearly  a  half-hour 
before  being  used.  Solution  V,  containing  1  c.c.  of  alkali,  was  used 
about  4  hours  after  being  prepared. 

Similar  experiments  were  made  with  solutions  of  phenolphthalein, 
partially  converted  into  the  quinoid  salt,  and  it  was  found  that  such 
solutions  are  fairly  stable.  Very  little  change  in  transparency  could  be 
detected  during  the  first  4  or  5  hours  after  the  solutions  were  prepared. 


TABLE  38. 
[I/Io  for  depth  of  solution  =20  mm.] 


X  =A.  U. 

I. 

II. 

III. 

IV. 

V. 

5773 

10.5 

10.1 

12.1 

10.1 

24.9 

5798 

15.2 

15.4 

16.5 

15.2 

31.0 

5823 

20.8 

20.2 

22.6 

21.4 

37.8 

5848 

28.7 

27.4 

30.3 

28.4 

45.6 

5947 

59.4 

57.7 

61.2 

60.3 

71.1 

RESULTS  WITH  PHENOLPHTHALEIN. 

The  six  solutions  recorded  below  were  prepared  from  the  following 
mother  solutions.  The  concentration  of  the  phenolphthalein  was 
6.697  X  10~5  gram-molecules  per  liter;  that  of  the  ammonium  hydroxide 
0.1448  gram-molecules  per  liter,  and  that  of  the  ammonium  chloride 
0.1  gram-molecules  per  liter. 

SOLUTIONS. 

I.  75  c.c.  phenolphthalein,  diluted  to  100  c.c. 

II.  75  c.c.  phenolphthalein,  2  c.c.  NH4OH,  1  c.c.  NH4C1,  diluted  to  100  c.c. 

III.  75  c.c.  phenolphthalein,  2  c.c.  NH4OH,  2  c.c.  NH4C1,  diluted  to  100  c.c. 

IV.  75  c.c.  phenolphthalein,  2  c.c.  NH4OH,  5  c.c.  NH4C1,  diluted  to  100  e.c. 
V.  75  c.e.  phenolphthalein,  2  c.c.  NH4OH,  10  c.c.  NH4C1,  diluted  to  100  c.c. 

VI.  75  c.c.  phenolphthalein,  1  c.c.  N  NaOH,  diluted  to  100  c.c. 


TABLE  39. 
i'1/Io  for  depth  of  solution  =20  mm.     Temperature  =  20°.] 


X  -A.  U. 

I. 

II. 

III. 

IV. 

V. 

VI. 

5773 

100.0 

30.8 

46.3 

70.7 

10.1 

5798 

100.0  !  37.8 

53.4 

76.3 

86.7 

14.3 

5823 

100.0 

45.3 

58.9 

78.5 

89.3 

20.2 

5848 

100.0 

51.7 

66.2 

83.0 

91.3 

27.4 

5948 

100.0 

83.7 

57.7 

Radiometric  Measurements  of  Constants  of  Indicators. 


67 


The  percentage  transmissions  for  these  solutions  are  given  in  table 
39.  The  concentrations  of  the  two  tautomeric  forms,  and  the  hydroly- 
sis and  ionization  constants  calculated  from  table  39  are  recorded  in 
table  40. 

TABLE  40. 


Solution. 

X  =A.  U. 

Lactoid 
form 
cXIO5. 

Quinokl 
salt 
cXIO6. 

OH  X  10s. 

KK/K, 
X105. 

K«X10">. 

Average 
K(X10'». 

II 

5773 

2.44 

2.58 

5.039 

4.77 

.70 

5798 

2.51 

2.51 

5.039 

5.05 

.60 

1.63 

5823 

2.54 

2.48 

5.039 

5.17 

.57 

5848 

2.46 

2.56 

5.039 

4.85 

.67 

III 

5773 

3.33 

1.69 

2.519 

4.96 

.64 

5798 

3.40 

1.62 

2.519 

5.27 

.54 

5823 

3.35 

1.67 

2.519 

5.06 

.60 

1.56 

5848 

3.42 

1.60 

2.519 

5.38 

.51 

5948 

3.40 

1.62 

2.519 

5.28 

1.53 

IV 

5773 

4.26 

0.757 

1.007 

5.68 

1.43 

5798 

4.32 

0.696 

1.007 

6.27 

1.20 

1.38 

5823 

4.26 

0.761 

1.007 

5.63 

1.44 

5848 

4.30 

0.720 

1.007 

6.02 

1.35 

V 

5798 

4.65 

0.369 

0.503 

6.37 

1.27 

5823 

4.67 

0.356 

0.503 

6.63 

1.22 

1.23 

5848 

4.67 

0.352 

0.503 

6.68 

1.21 

The  series  of  solutions  the  results  of  which  are  recorded  in  table  41 
were  prepared  from  another  solution  of  phenolphthlaein.  Its  concen- 
tration was  4.603  X  10~5  gram-molecules  per  liter.  The  concentrations 
of  the  ammonium  hydroxide  and  ammonium  chloride  solutions  were  the 
same  as  for  the  previous  series.  It  will  be  noticed  that  1  c.c.  of  a  solu- 
tion of  ammonium  hydroxide  was  used  instead  of  2  c.c.  as  in  the  former 
case. 

SOLUTIONS. 

I.  75  c.o.  phenolphthalcin,  1  c.c.  NH4OH,  0.5  c.c.  NH4C1,  diluted  to  100  c.c. 
II.  75  c.c.  phenolphthalein,  1  c.c.  NH4OH,  1  c.c.  NH4C1,  diluted  to  100  c.c. 

III.  75  c.c.  phenolphthalein,  1  c.c.  HN4OH,  2  c.c.  NH4C1,  diluted  to  100  c.c. 

IV.  75  c.c.  phenolphthalein,  1  c.c.  NH4OH,  2.5  c.c.  NH4C1,  diluted  to  100  c.c. 
V.  75  c.c.  phenolphthalein,  0.8  c.c.  N  NaOH,  diluted  to  100  c.c. 

TABLE  41. 
[I/Io  for  depth  of  solution  =20  mm.     Temperature  =20°.] 


X  =A.  U. 

I. 

II. 

III. 

IV. 

V. 

5773 

48.2 

62.2 

73.0 

81.3 

15.7 

5798 

54.2 

66.3 

78.7 

83.0 

22.7 

5823 

59.6 

70.7 

83.9 

86.0 

30.5 

5848 

66.4 

74.6 

85.1 

89.0 

37.9 

5948 

84.8 

87.7 

94.8 

62.2 

68 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


TABLE  42. 


Solution. 

X  =A.  U. 

Lactoid 
form 
cXIO5. 

Quinoid 
salt 
cXIO5. 

OHX106. 

Ktt./K« 
X105. 

K,X1010. 

Average 
KjXIO10. 

I 

5773 

2.09 

1.36 

5.039 

7.79 

1.04 

5798 

2.03 

1.42 

5.039 

7.24 

1.12 

1.14 

5823 

1.95 

1.50 

5.039 

6.56 

1.23 

5848 

1.99 

1.46 

5.039 

6.92 

1.17 

II 

5773 

2.56 

0.884 

2.519 

7.31 

1.11 

5798 

2.49 

0.953 

2.519 

6.61 

1.23 

5823 

2.44 

1.006 

2.519 

6.13 

1.32 

1.25 

5848 

2.41 

1.040 

2.519 

5.83 

.39 

5948 

2.49 

0.951 

2.519 

6.61 

.23 

III 

5773 

2.86 

0.587 

1.259 

6.15 

.32 

5798 

2.89 

0.557 

1.259 

6.53 

.24 

1.21 

5823 

2.94 

0.508 

1.259 

7.96 

.02 

5848 

2.87 

0.573 

1.259 

6.37 

.27 

IV 

5773 

3.05 

0.386 

1.007 

8.03 

.01 

5798 

3.01 

0.433 

1.007 

7.01 

1.15 

5823 

3.01 

0.436 

1.007 

7.01 

1.15 

1.00 

5848 

3.03 

0.413 

1.007 

7.42 

1.09 

5948 

3.06 

0.385 

1.007 

8.04 

1.01 

The  concentrations  of  the  mother  solutions  used  in  the  preparation 
of  the  solutions  for  the  following  series  of  measurements  were  the  same 
as  those  employed  for  table  39.  The  results  are  recorded  in  tables  43 
and  44. 

SOLUTIONS. 

I.  75  c.c.  phenolphthalein,  0.5  c.c.  NH4OH,  0.5  c.c.  NH4C1,  diluted  to  100  c.c. 
II.  75  c.c.  phenolphthalein,  0.5  c.c.  NH4OH,  1  c.c.  NH4C1,  diluted  to  100  c.c. 

III.  75  c.c.  phenolphthalein,  0.5  c.c.  NH4OH,  2  c.c.  NH4Cl,  diluted  to  100  c.c. 

IV.  75  c.c.  phenolphthalein,  0.8  c.c.  N  NaOH,  diluted  to  100  c.c. 

TABLE  43. 
[I/Io  for  depth  of  solution  =20  mm.     Temperature  =20°.| 


X=A.  U. 

I. 

II. 

III 

IV. 

5773 

63.2 

76.2 

87.3 

9.9 

5798 

65.8 

80.3 

89.2 

14.7 

5823 

72.0 

82.8 

91.5 

20.4 

5848 

74.3 

27.4 

5948 

88.8 

93.3 

57.6 

TABLE  44. 


Solution. 

X=A.  U. 

Lactoid 
form 
cXIO6. 

Quinoid 
salt 
X106. 

OHX105. 

KjKf 
X10S. 

KjXIO10. 

Average 
K4X10">. 

I 

5773 

4.026 

0.99 

2.519 

10.16 

0.80 

5798 

3.935 

1.09 

2.519 

9.07 

0.99 

5823 

3.985 

1.04 

2.519 

9.65 

0.83 

0.89 

5848 

3.875 

1.15 

2.519 

8.47 

0.97 

5948 

3.955 

1.07 

2.519 

9.28 

0.87 

II 

5773 

4.43 

0.59 

1.26 

9.46 

0.86 

5798 

4.45 

0.57 

1.26 

9.77 

0.83 

0^87 

5823 

4.42 

0.60 

1.26 

9.24 

0.88 

5948 

4.39 

0.62 

1.26 

8.82 

0.92 

III 

5773 

4.73 

0.30 

0.63 

10.07 

0.81 

5798 

4.73 

0.30 

0.63 

10.00 

0.81 

0.79 

5823 

4.74 

0.28 

0.63 

10.65 

0.75 

: 

Radiometric  Measurements  of  Constants  of  Indicators.  69 

The  concentration  of  phenolphthalein  is  not  a  factor  in  the  variations 
of  the  constants  in  the  different  tables.  The  following  shows  that  it  is 
possible  to  determine  Kj  without  knowing  the  concentration  of  the 
indicator. 

If  we  consider  equation  21 

LH  X  OH       Kg 
Q+          =  K, 

and  represent  the  total  concentration  of  phenolphthalein  by  T,  and  that 
of  the  quinoid  salt  by  Q,  we  have  from  equation  24 

In  (I/Ip)  X  T 
In  (I/Io)'  M 

LH  is  then  equal  to  (T  —  Q)  or 

,T  T       In  (I/Io)  X  T 

Q)  In  (I/Io)' 

Substituting  these  values  in  the  hydrolysis  equation,  we  have 

ln(I/Io)  X  T-l          - 

In  (I/Io)-     J'  Kw 


In  (I/Ip)  X  T  =  K, 

In  (I/Io)' 

Simplifying,  T,  the  concentration  of  the  phenolphthalein  disappears, 
and  we  have 

OH[ln  (I/Ip)'  -  In  (I/I0)]       K.  . 

In  (I/Io)  =  K, 

K,,  calculated  from  equation  29,  gives  the  same  value  as  when  calcu- 
lated from  the  equations  previously  derived. 

Since  it  is  not  necessary  to  know  the  concentration  of  the  indicator  to 
determine  its  ionization  constant,  the  variation  of  Kt,  shown  by  the 
different  tables,  must  be  due  to  other  causes.  It  will  be  noticed  from 
the  separate  tables  that  K;  decreases  with  the  hydroxyl  ion  concentra- 
tion. This  is  in  accordance  with  the  results  of  previous  investigators. 
When  the  hydroxyl  ion  concentration,  or  the  ratio  of  ammonium 
chloride  to  ammonium  hydroxide  remains  the  same,  K;  shows  a  marked 
decrease  with  decreasing  amounts  of  neutral  salts.  If  we  consider 
solutions  III  and  I  in  tables  40  and  44,  respectively,  it  will  be  seen  that 
decreasing  the  concentration  of  the  neutral  salt  four  times  changes  K, 
from  1.6X10"10  to  0.9X10~10.  This  is  in  accord  with  the  results  of 
Rosenstein,  who  has  made  a  very  careful  study  of  the  effect  of  neutral 
salts  on  the  ionization  constant  of  phenolphthalein.  It  will,  therefore, 
be  seen  that  the  discrepancies  in  the  value  of  K,  are  due  in  part  at  least 


70  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

to  the  effect  of  the  neutral  salt.  From  this  it  might  be  concluded  that 
K;  would  be  largest  for  those  solutions  containing  the  greatest  amount 
of  neutral  salt.  A  study  of  the  separate  tables,  table  40  for  instance, 
shows  that  just  the  reverse  is  true.  The  solution  that  contains  the 
largest  amount  of  neutral  salt  always  gives  the  smallest  concentration 
of  hydroxyl  ions,  and  the  K,  value  for  such  a  solution  is  much  lower 
than  for  a  solution  containing  less  neutral  salt  but  a  greater  concen- 
tration of  hydroxyl  ions.  It  is  thus  evident  that  K,  varies  not  only 
with  the  concentration  of  the  neutral  salt,  but  also  with  the  concen- 
tration of  the  hydroxyl  ions.  From  the  above  it  will  be  seen  that  if 
the  neutral  salt  had  no  effect,  the  variation  of  K;  in  the  separate  tables 
would  be  much  greater  than  it  actually  is.  It  is  therefore  obvious 
that  the  equilibrium  equations,  based  on  the  assumption  that  phenol- 
phthalein acts  as  a  monobasic  acid,  do  not  hold. 

Wegscheider  concludes  that  these  variations  may  be  accounted  for 
by  regarding  phenolphthalein  as  a  dibasic  acid.  The  results  of  Rosen- 
stein  make  it  appear  very  probable  that  this  indicator  does  act  as  a 
dibasic  acid.  The  limitations  of  this  method  and  the  presence  of 
neutral  salts  made  it  impossible  for  him  to  determine  satisfactorily  the 
ionization  constant  for  phenolphthalein,  and  the  question  must  there- 
fore be  regarded  as  still  open. 

The  radiometric  method  has  thus  far  been  applied  for  determining 
the  hydrolysis  and  ionization  constants  of  indicators.  Knowing  these 
values,  it  is  possible,  by  radiometric  means,  to  determine  from  them 
the  hydrolysis  constants  of  many  salts.  The  calculation  of  the  ioni- 
zation constants  of  weak  acids  and  bases  formed  by  the  hydrolysis  of 
these  salts  is  then  a  simple  matter. 

This  method  has  been  applied  in  a  preliminary  way  to  aluminum 
sulphate,  using  methyl  orange  as  the  indicator.  It  is  assumed  that  the 
secondary  and  tertiary  hydrolysis  of  this  salt  can  be  neglected,  and, 
on  this  assumption,  the  ionization  constant  of  the  base  formed  by  the 
primary  hydrolysis  has  been  calculated. 

The  calculations  of  the  constants  are  based  on  two  fundamental 
equilibrium  equations,  the  symbols  representing  gram-ions  per  liter. 
Expressing  the  equilibrium  relation  for  the  primary  hydrolysis  of 
aluminum  sulphate,  we  have 

A10H  X  H       Kw 

'  =  v~  (30) 


Al 
In  the  hydrolysis  equation  for  methyl  orange, 

AzOH  X  H       K,,, 
Q+  =  K, 


Radiometric  Measurements  of  Constants  of  Indicators. 


71 


K,  has  been  determined  and  Ku,  is  known.     The  concentration  of  the 

+ 

quinoid  ions  Q  is  given  by 


T[ln  (I/Ip)  -  In  (I/I0)'] 
In  (I/Io)"  -  In  (I/I,)' 


(3D 


as  explained  in  equation  19.     If  the  total  amount  of  indicator  is  T, 

+ 

AzOH  =  T  —  Q.     We  then  have  all  the  data  necessary  for  the  calcu- 

lation of  the  hydrogen  ion  concentration  in  equation  30.     The  con- 
+  +  +        + 

centration  of  the  ion  A1OH  in  the  same  equation  is  given  by  H  -f  Q, 

both  of  which  have  been  determined.     If  T'  represents  the  total  salt 


added  and  a  its  dissociation,  Al  is  (T'  —  A10H)  a. 

The  following  solutions  were  prepared  from  mother  solutions  of 
methyl  orange  and  aluminium  sulphate,  the  concentrations  being 
1.9865X  10~4  and  0.391  gram-molecules  per  liter,  respectively. 


SOLUTIONS. 

I.  60  c.c.  methyl  orange,  diluted  to  100  c.c. 
II.  50  c.c.  methyl  orange,  5  c.c.  Al2(SOi)s,  diluted  to  100  c.c. 

III.  50  c.c.  methyl  orange,  7  c.c.  A12(SO4)3,  diluted  to  100  c.c. 

IV.  SO  c.c.  methyl  orange,  10  c.c.  A12(SO4)8,  diluted  to  100  e.c. 

V.  50  c.c.  methyl  orange,  0.5  c.c.  cone.  i^SO*,  diluted  to  100  c.c. 


The  percentage  transmissions  given  by  these  solutions  are  recorded 
in  table  45. 

TABLE  45. 
[I/Io  for  depth  of  solution  =20  mm.] 


X=A.  U. 

I. 

II. 

III. 

IV. 

V. 

5698 

87.7 

47.7 

44.7 

39.7 

16.0 

5723 

88.9 

56.7 

53.1 

20.1 

5748 

90.3 

63.0 

59.1 

56.7 

31.0 

5773 

91.6 

70.2 

66.8 

63.6 

39.9 

5797 

93.3 

74.7 

73.6 

68.6 

48.4 

5823 

94.5 

78.4 

78.1 

75.5 

55.1 

5847 

95.6 

81.6 

78.5 

62.2 

The  constants  calculated  from  these  percentage  transmissions  are 
given  in  table  46.  It  has  been  found1  that  the  dissociation  of  the 
aluminium  sulphate  in  solutions  II,  III,  and  IV  is  approximately  35, 
32,  and  30  per  cent,  respectively. 

'Carnegie  Inst.  Wash.  Pub.  No.  170,  60  (1912). 


72 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 
TABLE  46. — Constants. 


Solution. 

X=A.  U 

+ 

Q 

10s. 

Average 
AzOH 
106. 

Average 
+ 
H 
104. 

Average 
+  + 
A1OH 
104. 

Average 
+  +  + 
Al 
103. 

Average 
KKK6 
X106. 

Average 
KSX10». 

II 

5698 

3.56 

5723 

3.01 

5748 

3.35 

5773 
5797 

3.20 
3.37 

6.56 

1.97 

2.307 

6.76 

6.75 

1.2 

5823 

3.44 

5847 

3.65 

III 

5698 

3.94 

5723 

3.44 

5748 
5773 

3.94 
3.79 

6.24 

2.24 

2.609 

8.66 

6.75 

1.2 

5797 

3.50 

5823 

3.51 

IV 

5698 

4.62 

5748 

4.33 

5773 
5797 

4.36 
4.66 

5.49 

3.12 

3.564 

11.62 

9.70 

0.84 

5823 

4.16 

5847 

4.54 

SUMMARY. 

A  satisfactory  radiomicrometer,  having  a  half -period  of  10  seconds 
and  a  sensibility  of  5  per  square  millimeter  of  exposed  vane  (candle  and 
scale  being  at  a  meter's  distance)  was  constructed.  Making  use  of  the 
radiomicrometer  and  the  grating  spectroscope,  a  radiometric  method 
was  worked  out  for  the  determination  of  the  ionization  constants  of 
indicators.  This  method  is  freer  from  objections  and  limitations  than 
any  method  previously  used.  It  serves  as  well  for  a  two-colored 
indicator  as  for  a  one-colored  indicator. 

Very  small  concentrations  of  colored  components  were  determined, 
and  it  has  been  shown  that  minute  concentrations  of  hydrogen  and 
hydroxyl  ions  can  be  quickly  and  accurately  estimated  by  means  of 
radiometric  measurements. 

Satisfactory  constants  were  obtained  for  the  ionization  of  methyl 
orange  as  a  base.  The  value  found  is  2.1  X  10~n. 

The  ionization  and  hydrolysis  constants  for  phenolphthalein  con- 
sidered as  a  monobasic  acid  are  far  from  being  satisfactory. 

From  the  known  ionization  constant  of  methyl  orange  and  from 
radiometric  measurements,  the  ionization  constant  of  a  very  weak 
base  and  the  hydrolysis  constant  of  one  of  its  salts  have  been  roughly 
determined.  The  method  can  likewise  be  applied  to  the  determina- 
tion of  the  ionization  constants  of  very  weak  acids,  and  the  hydrolysis 
constants  of  the  salts  formed  by  these  acids.  Work  is  now  in  progress 
in  this  laboratory  on  other  indicators  from  this  same  standpoint. 


CHAPTER  IV. 

RADIOMETRIC  MEASUREMENTS  OF  THE  IONIZATION  CONSTANTS  OF 

INDICATORS.-II. 


By  M.  G.  PATILUS  AND  J.  F.  HUTCHINSON. 


An  investigation  of  the  ionization  constants  of  methyl  orange  and 
phenolphthalein  has  already  been  published  by  Shaeffer,  Paulus,  and 
Jones.1  A  new  method,  based  upon  the  absorption  of  light  by  solutions 
of  indicators,  was  developed  for  the  determination  of  the  constants  of 
indicators,  and  is  now  presented.  It  was  shown  that  this  method 
serves  as  well  for  a  two-colored  as  for  a  one-colored  indicator.  The 
work  recorded  herein  is  to  be  regarded  as  a  continuation  of  the  original 
investigation,  and  the  purpose  is  to  test  the  applicability  of  the  method 
to  the  determination  of  the  ionization  constant  of  rosolic  acid.  A 
description  of  the  apparatus  used  has  already  been  given  in  detail  in 
the  original  paper. 

THEORETICAL  DISCUSSION. 

Considering  first  of  all  that  rosolic  acid  is  monobasic,2  the  ionization 
constant  K,  is  expressed  by  the  simple  equilibrium  equation 

(H+)  X  (15) 
(HIn) 

If,  then,  the  hydrogen  ion  concentration  of  the  indicator  solution  is 
fixed,  the  ratio  (In)  /(HIn)  at  equilibrium  can  be  determined.  It 
has  been  shown  that  the  percentage  transmission  of  a  solution,  such 
as  that  of  rosolic  acid,  containing  two  absorbing  components  is  given 
by  the  equation 

=  -Kc  -  K'd  (2) 


where  c  and  c\  are  the  concentrations  of  the  two  absorbing  components, 
and  K  and  K'  are  constants  depending  upon  the  nature  of  the  absorbing 
components  and  the  wave-length  of  light  employed.  Applying  this 
equation  to  rosolic  acid,  let  c  represent  the  concentration  of  the  red 
component,  or  (In)  in  equation  1,  and  let  c\  represent  the  concentration 
of  the  yellow  component  or  (HIn).  Equation  2,  then,  will  represent 
the  percentage  transmission  for  some  given  depth  of  an  incompletely 
transformed  solution  of  rosolic  acid.  In  a  solution  containing  a  large 
excess  of  acid  c  =  0,  and  equation  2  becomes 

In  (I/I0)'=  -K'Cl=-KT  _  (3) 

»Journ.  Amer.  Chem.  Soc.,  37,  776  (1915). 

The  behavior  of  rosolic  acid  as  a  dibasic  acid  will  be  discussed  later. 

73 


74  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

If  a  large  excess  of  alkali1  is  added  to  the  indicator  solution  Ci  =  0,  and 
equation  2  reduces  to 

ln(I/I0)"=-Kc  =   -KT  (4) 

where  T  equals  the  total  concentration  of  the  indicator  in  solution.  If 
the  percentage  transmissions  are  determined  for  the  same  depth  of 
solution  and  for  the  same  wave-length  of  light,  and  if  the  total  concen- 
tration of  the  indicator  is  the  same  in  all  solutions,  then  the  values  of 
the  constants  K  and  K'  given  by  equations  3  and  4  can  be  substituted 
in  equation  2,  whence 

TXln  (I/I0)  =  ln  (I/I0)"Xc  +ln  (I/I0)'X  c,  (5) 

Since  the  total  concentration  of  the  indicator  T  is  always  equal  to  the 
sum  of  the  two  components,  T  =  c  -\-c\,  equation  5  reduces  to 

In  (I/I,)'  -  In  (I/Ip)  (6) 

:  -" 


The  ratio  c/Ci  =  (In)/HIn)  can  be  determined  from  equation  6. 
(I  /I0)  is  the  percentage  transmission  for  some  given  depth  of  the 
incompletely  transformed  solution  under  investigation,  for  some  wave- 
length of  light;  (I  /Io)'  the  percentage  transmission  of  the  indicator 
solution  completely  transformed  into  the  yellow  component  for  the 
same  wave-length;  and  (I/Io)"  the  percentage  transmission  for  the 
same  depth  of  indicator  solution  completely  transformed  into  the  red 
component,  for  the  same  wave-length  of  light.  The  total  concentra- 
tion of  the  indicator  in  these  three  solutions  must,  of  course,  be  the 
same,  but  the  total  concentration  need  not  be  known. 

Returning  now  to  equation  1,  the  method  of  obtaining  all  the  data 
necessary  for  calculating  the  ionization  constant  is  known,  except  that 
for  determining  the  concentration  of  the  hydrogen  ion.  This  was  fixed 
by  solutions  of  disodium  phosphate  containing  varying  amounts  of 
hydrochloric  acid.  The  addition  of  hydrochloric  acid  converts  the 
hydrophosphate  ion  almost  quantitatively  into  the  dihydrophosphate 
ion.  The  hydrogen  ion  concentration  is  given  in  such  a  solution  by2 

H+  -  1-95  X  IP"7  (H2P04)  (7) 

(HP04) 

If  we  represent  by  a  the  concentration  of  the  added  hydrochloric  acid, 
and  by  b  the  concentration  of  the  disodium  phosphate,  then  equation  7 
becomes 

1.95  X  IP"7  X  otti 

(b  -  a)  a2 

where  0.1  and  a2  represent  respectively  the  dissociations  of  the  mono- 
and  disodium  phosphates  present  at  equilibrium. 

*By  "a  large  excess"  is  meant  sufficient  alkali  to  convert  the  indicator  entirely  into  its  red 
component. 

'The  value  of  the  constant  was  taken  from  the  work  of  Abbott  and  Bray:  Journ.  Amer.  Chem. 
Soc.,  31,  760  (1909). 


Radiometric  Measurements  of  Constants  of  Indicators.  75 

If  the  quantity  of  disodium  phosphate  in  the  solutions  investigated 
is  always  kept  the  same,  the  hydrogen  ion  concentration  can  be  varied 
simply  by  the  addition  of  different  amounts  of  hydrochloric  acid.  In 
this  case  the  total  salt  concentration  is  constant.  This  is  extremely 
desirable,  as  it  has  been  shown  by  Rosenstein1  that  neutral  salts  have 
a  great  effect  upon  the  fraction  of  the  indicator  transformed.  The 
value  of  the  ionization  constant  in  the  case  of  phenolphthalein  is 
doubled  by  increasing  the  total  salt  concentration  from  0.03  to  0.40 
normal. 

PRELIMINARY  WORK  ON  ROSOLIC  ACID. 

Three  stock  solutions  were  prepared,  all  solutions  being  made  up 
at  20°  with  conductivity  water.  The  stock  solution  of  disodium  phos- 
phate was  prepared  from  a  pure  sample  obtained  from  Kahlbaum.  Its 
concentration,  0.1036  gram-molecules  per  liter,  was  determined  gravi- 
metrically  as  magnesium  pyrophosphate.  The  concentration  of  the 
stock  solution  of  hydrochloric  acid  was  0.08085  normal.  The  stock 
indicator  solution  was  prepared  by  dissolving  about  0.4  gm.  of  an 
excellent  sample  of  rosolic  acid  obtained  from  Merck  in  2  liters  of 
conductivity  water.  The  total  quantity  of  the  indicator  did  not 
dissolve,  but,  as  has  been  shown,  it  is  not  necessary  to  know  the  concen- 
tration of  the  indicator  employed. 

The  incompletely  transformed  solutions  to  be  tested  were  prepared 
from  the  stock  solutions,  so  that  all  contained  the  same  amounts  of 
indicator  and  disodium  phosphate,  but  different  amounts  of  hydro- 
chloric acid.  This  procedure  was  followed  to  keep  the  total  salt  con- 
centration the  same  in  all  solutions.  The  volume  of  each  solution 
was  100  c.c.  The  percentage  transmissions  (I/Id),  were  taken  with 
a  20  mm.  depth  of  each  solution,  and  for  the  same  5  wave-lengths  of 
light.  As  explained  in  the  original  article,  the  percentage  transmissions 
were  determined  by  a  differential  method,  which  avoided  the  necessity 
of  introducing  certain  correction  factors  due  to  the  glass  ends  with 
which  the  cells  were  provided. 

A  consideration  of  equation  6  will  show  that  the  method  of  calcu- 
lating the  ratio  of  the  red  to  the  yellow  component  of  any  incompletely 
transformed  indicator  solution,  not  only  depends  upon  the  transmission 
of  the  solution  under  investigation,  but  also  upon  the  transmission  of 
an  indicator  solution  containing  a  large  excess  of  acid,  in  which  the 
indicator  is  totally  transformed  into  its  yellow  constituent;  and  also 
upon  the  transmission  of  an  indicator  solution  having  a  large  excess 
of  alkali  in  which  the  indicator  is  totally  transformed  into  its  red  com- 
ponent. Table  47  gives  the  results  of  a  series  of  measurements  made 
upon  the  indicator  solution  which  has  an  excess  of  alkali.  All  solu- 
tions contain  50  c.c.  of  the  stock  solution  of  rosolic  acid  plus  the 

'Abbott  and  Bray:  Journ  Amer.  Chem.  Soc.,  34,  1128  (1912).     See  also:  Ibid.,  37,  804  (1915) . 


76 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


amount  of  N/l  NaOH  indicated  in  the  table,  the  solutions  being  in  all 
cases  diluted  to  100  c.c.  The  percentage  transmissions  are  given  for 
5  wave-lengths  of  light  between  A  =  0.56M  and  X  =  0.58/1,  which  is  the 
region  of  the  spectrum  employed  throughout  this  investigation.  In 
certain  cases  the  results  of  duplicate  measurements  are  given,  which 
indicate  in  a  general  way  the  accuracy  of  the  results. 

TABLE  47. 
[I/Io— depth  of  solution  =  20  mm.] 


"0 

0.5  c.c.  N/l  NaOH  after— 

w  e 

u  O  -C 

6£3 

X  =A.  U. 

°    rt      lH 

5  min. 

10  min. 

20  min. 

30  min. 

55  min. 

5  hrs. 

24  hrs. 

ro  ""-^.2 

5598 

22.6 

16.1 

12.5 

10.7 

9.08 

7.55 

6.25     6.37 

6.53     6.53 

9.16 

5648 

36.1 

30.6 

27.4 

22.6 

20.0 

20.0 

17.0     16.7 

17.6     16.7 

19.1 

5698 

52.3 

44.8 

44.2 

38.8 

37.9 

37.0 

34.5     34.5 

32.2     32.2 

35.6 

5748 

64.8 

60.6 

58.3 

55.6 

52.8 

51.4 

52.4     50.7 

49.3     48.7 

50.7 

5798 

77.0 

73.3 

71.1 

68.4 

64.1 

64.8 

66.6     66.2 

63.5     63.7 

66.3 

Column  2,  table  47,  gives  the  percentage  transmissions  of  an  indicator 
solution  containing  0.5  c.c.  N/l  NaOH,  the  transmissions  being  deter- 
mined within  a  short  time  after  the  solution  was  prepared.  Columns 
3  to  7,  inclusive,  give  the  percentage  transmissions  after  the  same 
solution  had  stood  for  various  intervals  of  time  up  to  55  minutes. 
Columns  8  and  9  give  the  percentage  transmissions  of  new  indicator 
solutions  containing  the  same  amount  of  indicator  and  alkali,  after  these 
solutions  had  stood  for  5  and  24  hours  respectively.  It  will  be  observed 
that  the  percentage  transmissions  for  any  given  wave-length  of  light 
become  constant  after  the  solutions  have  stood  between  1  and  5  hours. 

The  fact  that  solutions  of  rosolic  acid  containing  an  excess  of  alkali 
become,  on  standing,  less  and  less  transparent  to  yellow  light,  clearly 
indicates  that  the  concentration  of  the  red  component  (In)  present  is 
becoming  greater  and  greater,  with  a  resulting  decrease  in  the  con- 
centration of  the  yellow  component  (HIn);  since  the  greater  the 
concentration  of  the  red  component  the  more  opaque  the  solution 
becomes  to  yellow  light. 

According  to  the  most  recent  views1  concerning  the  cause  of  color 
production  by  indicators  of  the  aurine  type,  the  color  is  not  due  simply 
to  the  presence  of  a  quinoid  group  as  such,  but  to  an  inter-  or  intra- 
molecular combination  of  the  metallic  phenolate  with  the  quinoid 
complex.  It  is  very  probably  true,  in  the  case  of  rosolic  acid,  that  this 
combination  between  the  metallic  phenolate  and  the  quinoid  complex 
takes  place  rather  slowly,  with  a  corresponding  intensification  of  the 
red  color. 

It  will  be  observed  that  the  transmission  values  recorded  in  column 
10,  table  47,  of  an  indicator  solution  containing  3  c.c.  N/lNaOH  are 

"Amer.  Chem.  Journ.,  39,  537.  650,  and  651  (1908). 


Radiometric  Measurements  of  Constants  of  Indicators. 


77 


considerably  higher  than  those  recorded  in  columns  8  and  9,  for  a  solu- 
tion containing  0.5  c.c.  N/lNaOH.  In  both  cases  the  solutions  have 
come  to  equilibrium,  since  it  has  been  shown  that  equilibrium  is  estab- 
lished after  the  solutions  have  stood  between  1  and  5  hours.  When 
very  much  larger  amounts  of  alkali  are  added  (say  10  c.c.),  a  very 
perceptible  bleaching  takes  place.  In  solutions,  then,  containing  an 
excess  of  sodium  hydroxide,  two  opposing  reactions  take  place;  first,  a 
gradual  intensification  of  the  red  color  brought  about  very  probably 
by  a  time  reaction  between  the  metallic  phenolate  and  the  quinoid 
complex;  and  second,  a  bleaching  of  the  red  color  which  is  greater 
the  larger  the  amount  of  the  alkali  added.  Since  it  is  necessary  to 
know  the  true  percentage  transmission  (I/I0)"  of  a  solution  completely 
transformed  into  the  red  component,  in  order  to  determine  the 
ratio  c/Ci,  it  follows  that  the  bleaching  must  be  avoided.  This  result 
can  be  obtained  by  keeping  the  concentration  of  the  alkali  as  small 
as  possible,  only  adding  sufficient  to  transform  completely  the  indi- 
cator. The  obvious  method  was  to  increase  gradually  the  hydroxyl 
ion  concentration  until  the  solutions  were  shown  to  be  completely 
transformed.  Table  48  gives  the  results  of  such  a  determination.  All 
solutions  contain  50  c.c.  of  the  stock  solution  of  indicator  plus  the 
amount  of  disodium  phosphate  and  hydrochloric  acid  indicated  in 
the  table.  All  solutions  were  allowed  to  stand  a  sufficient  length  of 
time  for  equilibrium  to  be  established. 


TABLE  48. 
[(I/Io)"—  depth  of  solution 


20  mm.] 


25  c.o.Na2HPO4 

25  c.c.  NasHPO4 

25  c.c.  NajHPO4 

25  c.c.  Na2HPO4 

X=A.  U. 

1  c.c.  HC1. 

0.5  c.c.  HC1. 

0  c.c.  HC1. 

0  c.c.  HC1. 

After  24  hrs. 

After  48  hrs. 

After  48  hrs. 

After  96  hrs. 

5598 

4.44         4.35 

4.68 

4.68         4.76 

5.00 

5648 

13.7         14.0 

11.1 

10.8         11.0 

9.6 

5698 

22.8         22.8 

19.4 

19.4         18.8 

19.0 

5748 

40.0         40.0 

33.3 

34.1         34.9 

34.5 

5798 

51.6         51.0 

48.4 

48.4         49.4 

48.8 

Solution  1,  table  48,  in  which  the  hydrogen  ion  concentration  is 
0.8266  X  10~8,  gives  higher  values  for  the  percentage  transmissions  than 
solutions  2  and  3,  in  which  the  hydrogen  ion  concentrations  are,  respec- 
tively, 0.4157  X10~8  and  0.043  XlO"8.  This  shows  that  solution  1  is 
not  completely  transformed,  since  an  increase  in  the  red  component 
and  a  corresponding  decrease  in  the  yellow  component  make  the  solu- 
tion more  opaque  to  yellow  light.  Solution  2  must,  however,  be  com- 
pletely transformed,  since,  when  the  hydrogen  ion  concentration  is  still 
further  decreased,  as  is  the  case  in  solution  3,  the  transmission  values 
remain  the  same.  The  values  of  the  percentage  transmissions  recorded 
in  columns  2,  3,  and  4  are  the  values  to  be  subsituted  for  (I/I0)"  in 
equation  6. 


78 


Condwtivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


In  order  to  ascertain  if  this  same  decrease  in  the  transmissions,  cor- 
responding to  an  increase  in  the  red  component,  would  take  place  on 
standing  with  solutions  of  indicators  which  are  incompletely  trans- 
formed, the  following  series  of  solutions  were  prepared: 


SOLUTIONS. 

No.  1.  50  c.c.  rosolic  acid;  25  c.c.  Na2HPO4; 

No.  2.  50  c.c.  rosolic  acid;  25  c.c. 

No.  3.  50  c.c.  rosolic  acid;  25  c.c. 

No.  4.  50  c.c.  rosolic  acid;  25  c.c. 

TABLE  49. 
[(I /Io)  — depth  of  solution 


1  c.c.  HC1  diluted  to  100  c.c. 

5  c.c.  HC1  diluted  to  100  c.c. 
10  c.c.  HC1  diluted  to  100  c.c. 
15  c.c.  HC1  diluted  to  100  c.c. 

=  20  mm.] 


Solution. 

X=A.  U. 

After  5  hrs. 

After  16  hrs. 

After  24  hrs. 

1 

5598 

16.7  16.3 

4  .  44  4  .  35 

5648 

32.7  30.2 

13.7  14.0 

5698 

42  .  8  43  .  2 

22  .  8  22  .  8 

5748 

59  .  1  57  .  4 

40.0  40.0 

5798 

69.7  70.3 

51.6  51.6 

2 

5598 

30.8  30.8 

25.7  25.0 

26.4  26.4 

5648 

39.3  39.6 

35.4  35.9 

37.5  35.7 

5698 

52.6  51.7 

48.3  40.0 

50.0  49.3 

5748 

66.1  64.5 

60.8  60.7 

60.3  59.4 

5798 

72.7  74.3 

71.0  71.0 

70.7  70.6 

3 

5598 

49.0  49  0 

50.7  50.7 

49.0  49.1 

5648 

57.8  57.8 

60.0  59.3 

60.0  58.8 

5698 

66.7  68.3 

68.7  68.7 

70.6  70.6 

5748 

73.8  73.4 

78.5  77.7 

76.2  77.4 

5798 

83.3  83.6 

84.2  84.2 

84.4  83.0 

4 

5598 

66.1  67.3 

68.0  67.2 

67.3  67.8 

5648 

72.4  74.1 

74.9  74.3 

76.4  76.4 

5698 

81.6  80.3 

81.3  80.4 

81.0  79.5 

5748 

83.2  83.2 

87.2  87.2 

85.5  85.5 

5798 

88.2  89.4 

90.2  90.2 

87.7  87.8 

In  table  49  are  given  the  percentage  transmissions  for  a  depth  equal 
to  20  mm.  of  the  above  series  of  solutions.  The  transmissions  were 
determined  after  the  solutions  had  stood  for  intervals  of  5, 16,  and  24 
hours.  The  transmission  values  recorded  in  columns  1  and  3  were 
determined  with  the  same  series  of  solutions.  A  new  series  of  solutions 
was  prepared  for  the  16-hour  determination,  the  results  of  which  are 
recorded  in  column  2.  Solution  No.  1  was  not  made  up  for  this  deter- 
mination. Duplicate  measurements  are  given  in  every  case. 

An  examination  of  the  transmission  values  recorded  for  solutions  1 
and  2,  table  49,  shows  that  a  decrease  in  the  transmissions  of  incom- 
pletely transformed  solutions  of  rosolic  acid  also  takes  place  on  standing. 
Solution  2,  in  which  the  hydrogen  ion  concentration  is  4.484  X1CT8, 
has  come  to  equilibrium  between  5  and  16  hours,  as  is  shown  by  the  fact 
that  the  transmission  values  become  constant  after  16  hours.  In 
solutions  3  and  4,  in  which  the  hydrogen  ion  concentrations  are  respec- 
tively 10.53  XlCT8  and  19.52  XlCT8,  equilibrium  was  established  before 
standing  5  hours.  This  points  to  the  conclusion  that  in  the  incom- 
pletely transformed  solutions  of  the  indicator  the  more  alkaline  the 
solution  the  greater  the  time  before  equilibrium  is  established. 


Radiometric  Measurements  of  Constants  of  Indicators. 


79 


RESULTS  WITH  ROSOLIC  ACID. 

For  the  first  determination  of  the  indicator  constant  K<  the  following 
series  of  solutions  was  used: 

SOLUTIONS. 

No.  1.  50  c.c.  rosolic  acid;  25  c.c.  NajHPO,;  5  c.c.  HC1,  diluted  to  100  c.c. 
No.  2.  50  c.c.  rosolic  acid;  25  c.c.  NajHPO^  10  c.c.  HC1,  diluted  to  100  c.c. 
No.  3.  50  e.c.  rosolic  acid;  25  c.c.  Na2HPO4;  15  c.c.  HC1.  diluted  to  100  c.c. 

These  solutions  were  prepared  from  the  stock  solutions,  the  concen- 
trations of  which  have  already  been  given.  All  solutions  were  allowed 
to  stand  for  24  hours,  which  time,  according  to  the  results  of  the  prelimi- 
nary work,  was  amply  sufficient  for  equilibrium  to  be  established.  In 
table  50  are  given  the  percentage  transmissions  for  a  depth  equal  to 
20  mm.  of  each  of  these  solutions,  the  values  being  in  every  case  the 
average  of  two  measurements.  These  are  the  values  to  be  substituted 
for  (I /Io)  in  equation  6.  As  has  previously  been  explained,  the  ratio 
c/Ci  also  depends  upon  the  percentage  transmission  (I/Io)"  of  an  indi- 
cator solution  completely  transformed  into  the  red  component,  and 
also  upon  the  percentage  transmission  (I/Io)'  of  an  indicator  solution 
completely  transformed  into  the  yellow  constituent.  The  transmissions 
(I/Io)"  used  are  the  averages  of  those  recorded  in  columns  2,  3,  and  4, 
table  48.  It  was  found  that  solutions  containing  50  c.c.  of  the  stock 
solution  of  the  indicator  plus  the  necessary  amount  of  hydrochloric 
acid  to  convert  the  indicator  entirely  into  the  yellow  component  were 
completely  transparent  to  the  5  wave-lengths  of  light  used.  The 
transmissions  (I/Io)'  are  therefore  in  every  case  100  per  cent. 

TABLE  50. 


First  solution. 

Second  solution. 

Third  solution. 

H+X108=4.484. 

H+X108  =  10.53. 

H+X108  =  19.52. 

o 

Average 

Average 

Average 

percentage 
trans- 

e/e, 

K,X 

10". 

percentage 
trans- 

c/c, 

K4X 
10s. 

percentage 

trans- 

c/c, 

K<X 
10s. 

missions 

missions 

missions 

5598 

26.4 

0.803 

3.60 

49.0 

0.306 

3.23 

67.5 

0.148 

2.89 

5648 

36.6 

0.811 

3.64 

59.4 

0.302 

3.19 

76.4 

0.136 

2.66 

5698 

49.6 

0.734 

3.29 

70.6 

0.266 

2.81 

80.2 

0.154 

3.05 

5748 

59.8 

0.920 

4.13 

76.8 

0.326 

3.44 

85.5 

0.171 

3.34 

5798 

70.6 

0.937 

4.20 

83.7 

0.329 

3.48 

87.7 

(0.223) 

(4.35) 

In  table  50  the  ratios  c/d  =  (In)/(HIn)  were  calculated  from  equa- 
tion 6  and  the  constants  K,  from  equation  1.  The  hydrogen  ion  con- 
centrations were  calculated  by  means  of  equation  7,  ai  and  a2  being 
interpolated  from  the  percentage  ionizations  of  monosodium  and 
disodium  phosphates  at  various  dilutions  given  by  Abbott  and  Bray.1 

Another  series  of  solutions  described  below  was  prepared  and 
allowed  to  stand  16  hours,  in  which  time  all  of  them  had  come  to 


'Amer.  Chem.  Journ.,  31,  729  (1909). 


80 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


equilibrium.  The  percentage  transmissions,1  calculated  ratios  c/Ci, 
and  the  constants  K;,  are  given  in  table  51.  The  same  values  for 
(I/ Io)'  and  (I/Io)"  were  used  as  for  the  preceding  determination. 


SOLUTIONS. 

No.  1.  50  c.c.  rosolic  acid:  25  c.c.  Na2HPO4 

No.  2.  50  c.c.  rosolic  acid;  25  c.c.  Na2HPO4 

No.  3.  50  c.c.  rosolic  acid;  25  c.c.  Na2HPO4 

No.  4.  50  c.c.  rosolic  acid;  25  c.c.  Na2HPO4 

TABLE  51. 


3  c.c.  HC1,  diluted  to  100  c.c. 

5  c.c.  HC1,  diluted  to  100  c.c. 
10  c.c.  HC1,  diluted  to  100  c.c. 
15  c.c.  HC1,  diluted  to  100  c.c. 


X  =A.  U. 

First  solution. 
H+X108=2.533. 

Second  solution. 
H+X108=4.484. 

Average 
percentage 
transmissions. 

C/Cl 

K,X108. 

Average 
percentage 
transmissions. 

efc 

K,X108. 

5598 
5648 
5698 
5748 
5798 

12.1 
20.8 
32.0 
48.1 
60.9 

2.28 
2.33 
2.20 
2.14 
2.22 

5.77 
5.89 
5.56 
5.42 
5.63 

25.4 
35.7 
48.3 
60.8 
71.0 

0.822 
0.847 
0.785 
0.849 
0.906 

3.69 
3.81 
3.53 
3.82 
4.07 

X  =A.  U. 

Third  solution. 
H+X10"=10.53. 

Fourth  solution. 
H+X108  =  19.52. 

Average 
percentage 
transmissions. 

c/c, 

K4X10». 

Average 
percentage 
transmissions. 

c/c, 

K,X10S. 

5598 
5648 
5698 
5748 
5798 

50.7 
59.7 
68.7 
78.1 
84.2 

0.288 
0.298 
0.293 
0.298 
0.314 

3.03 
3.14 
3.08 
3.14 
3.30 

67.6 
74.6 
80.8 
87.2 
(90.2) 

0.148 
0.150 
0.146 
0.146 
(0.167) 

2.89 
2.93 
2.89 
2.85 
(3.26) 

The  values  of  K,,  table  51,  show  a  steady  decrease  from  5.65  XlCT8 
to  2.89  X  10~8,  as  the  solution  becomes  less  alkaline.  The  average  con- 
stant is  3.91  X  10~8.<2)  The  constants  were  determined  on  the  assump- 
tion that  rosolic  acid  is  monobasic;  as  is  well  known,  the  indicator  is 
dibasic,  and  the  decrease  in  the  constants  with  decreasing  alkalinity  was 
expected.  Rosolic  acid  actually  dissociates  in  two  stages  according  to 
the  equations  H2In  =  HIn  +  H+  and  HIn  =  IS  -f  H+. 

BEHAVIOR  OF  ROSOLIC  ACID  AS  A  DIBASIC  ACID. 

As  to  the  two  ions  HIS  and  II,  three  assumptions  can  be  made:  (!) 
that  the  intermediate  ion  HIn  is  yellow  and  the  secondary  ion  In  is  red; 
(2)  that  the  ionJIIn  is  red  and  the  ion  In  is  yellow;  (3)  that  both  the 
ions  HIn  and  IH  are  red.  In  the  first  case  the  ratio  of  the  red  to  the 
yellow  component  will  be  given  by 

c/Cl  =  HI5  +  H2In  (9) 


'The  average  values  of  two  determinations  are  given. 

'The  value  given  by  Salm,  in  Zeit.  phys.  Chem.,  57,  496  (1907),  is  1.1  X10~8. 


Radiometric  Measurements  of  Constants  of  Indicators. 


81 


In  addition,  the  equilibrium  equations  given  below  are  to  be  considered  : 


(H+)  (HIE) 


=  K, 


(H+)  (In) 
(HIn) 


(H2In) 
By  combining  these  three  equations  we  obtain 

~K.~K. 

C/Ci    = 


H+(H+  + 


(10) 


(11) 


In  the  second  case,  when  the  ion  HIn  is  red  and  the  ion  In  is  yellow, 
the  ratio  becomes 

HIn 

(12) 


and  from  equation  10 


C/Ci  = 


H2In  +  In 
H+Kj 


H2 


(13) 


In  the  third  case,  i.  e.,  when  both  ions  HIn  and  In  are  red,  the  ratio  is 
given  by 

-       HE  +  IS  (14) 


whence,  c/Ci  = 


H2In 
H+  K!  + 


"> 


H+2 


(15) 


These  equations,  11,  13,  and  15,  were  tested  by  substituting  the 
experimental  values  of  the  ratios  c/Ci  and  the  hydrogen  ion  concen- 
trations given  for  solutions  1  and  4,  table  51,  in  the  equations,  and 
solving  for  the  constants  KI  and  K2.  The  constants  were  then  used  to 
calculate  the  ratios  for  solutions  2  and  3.  Table  52  gives  the  values 
of  the  ratios  c/Ci  calculated  from  the  various  equations,  and  also  the 
observed  experimental  ratios  which  are  the  averages  of  those  given  in 
table  51. 

TABLE  52. 


Values  of  c/c. 

Solution. 

H+X108. 

Observed. 

Calc.byll. 

Calc.  by  13. 

Calc.  by  15. 

1 

2.535 

2.23 

2.23 

2.23 

2.230 

2 

4.484 

0.842 

1.12 

0.953 

3 

10.53 

0.298 

0.371 

0.308 

4 

19.52 

0.148 

0.148 

0.148 

0.148 

On  the  first  assumption,  namely,  that  the  ion  HIn  is  yellow  and  the 
ion  In  is  red,  the  ratios  calculated  by  equation  11,  for  solutions  2  and  3, 
do  not  at  all  agree  with  the  experimentally  determined  values.  On  the 
assumption  that  the  intermediate  ion  HIn  is  red  and  the  secondary  ion 
In  is  yellow,  the  constant  K2,  equation  13,  was  found  to  be  a  negative 


'The  development  is  essentially  the  same  as  given  by  Rosenstein  for  phenolphthalein. 
Amor.  Chem.  Soc.,  36,  1124  (1912). 


Journ. 


82  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

quantity.  This,  in  itself,  proves  the  absurdity  of  the  assumption; 
consequently,  no  ratios  for  solutions  2  and  3_were  calculated.  On  the 
assumption,  however,  that  both  the  ions  HIn  and  In  are  red,  the  ratios 
calculated  by  equation  15  agree  fairly  closely  with  the  experimentally 
determined  values.  The  agreement  is  as  close  as  could  be  expected, 
considering  that  in  work  of  this  character  so  many  different  sources  of 
error  are  possible.  The  results  lead  to  the  conclusion,  therefore,  that 
rosolic  acid  acts  as  a  dibasic  acid,  and  furthermore,  show  that  both  the 
primary  and  secondary  ions  are  intensely  colored. 

In  the  case  of  phenolphthalein,  it  has  been  shown  by  Rosenstein  that 
the  colored  form  of  the  indicator  is  only  produced  in  appreciable  quan- 
tity where  the  second  hydrogen  of  the  indicator  acid  is  replaced  by  the 
base.  This  is  in  accordance  with  the  theory  already  referred  to,  that 
the  cause  of  the  color  production  is  due  to  a  combination  between 
the  metallic  phenolate  and  the  quinoid  complex,  since  it  is  only  where 
the  second  hydrogen  is  replaced  that  the  formation  of  the  quinoid 
phenolate  complex  is  possible.  In  the  case  of  rosolic  acid,  however, 
the  quinoid  phenolate  complex  can  be  found  when  the  first  hydrogen 
of  the  indicator  acid  is  replaced  by  the  base.  The  experimental  fact, 
therefore,  that  the  intensely  colored  form  of  the  indicator  is  produced 
when  the  first  hydrogen  is  replaced  by  the  base,  is  perfectly  in  accord 
with  the  theories  advanced. 

SUMMARY. 

1 .  An  intensification  of  the  red  color  of  solutions  of  rosolic  acid  incom- 
pletely transformed  by  the  addition  of  alkali,  was  found  to  take  place 
when  such  solutions  were  allowed  to  stand,  the  time  reaction  being  in 
all  probability  due  to  a  slow  union  of  the  metallic  phenolate  with  the 
quinoid  complex. 

2.  In  solutions  of  rosolic  acid  containing  a  large  excess  of  alkali  a 
perceptible  bleaching  of  the  red  color  was  also  indicated. 

3.  The  ratio  (c/Ci)  of  the  red  to  the  yellow  component  has  been 
determined  for  indicator  solutions  of  various  hydrogen  ion  concentra- 
tions, using  the  radiometric  method  developed  in  the  original  article. 

4.  The  values  of  the  ionization  constant  of  rosolic  acid  calculated 
from  the  ratio  c/c\,  on  the  assumption  that  the  indicator  acid  is  mono- 
basic, were  found  to  decrease  with  decreasing  alkalinity.     When  the 
hydrogen  ion  concentration  was  increased  from  2.533X10"8  to  19.52  X 
1(T8,  the  total  salt  concentration  being  0.0259  normal,  the  value  of  the 
ionization  constant  K,  was  found  to  decrease  from  5.65  X  10~8  to 
2.89  XlO~8. 

5.  It  was  found  that  this  variation  in  the  constants  could  be  ex- 
plained by  regarding  the  indicator  as  a  dibasic  acid,  and  it  was  further- 
more shown  that  the  intensely  colored  form  of  the  indicator  is  formed 
when  the  first  hydrogen  of  the  indicator  acid  is  replaced  by  the  base. 


CHAPTER  V. 

THE  ACTION  OF  SALTS  WITH  WATER  OF  HYDRATION  AND  WITHOUT 
WATER  OF  HYDRATION  ON  THE  VELOCITY  OF  SAPONIFICATION 
OF  ESTERS.  

BY  J.  E.  L.  HOLMES. 


Jones  and  Anderson,1  in  their  work  on  the  absorption  spectra  of 
solutions,  studied  the  absorption  spectra  of  neodymium  chloride  in 
water,  in  methyl  alcohol,  and  in  mixtures  of  these  two  solvents.  They 
found  two  sets  of  absorption  spectra  corresponding,  the  one  to  the 
aqueous  solution,  the  other  to  the  alcoholic.  In  the  mixture  of  these 
solvents,  both  of  these  spectra  wrere  obtained  when  the  water  was 
present  in  smaller  quantities  than  15  per  cent.  With  decrease  in  the 
percentage  of  water,  the  alcoholic  spectrum  increased  in  intensity. 
The  water  spectrum  and  the  alcohol  spectrum  were  found  to  be  quite 
different  from  one  another,  and  do  not  change  over  into  each  other 
when  the  composition  of  the  mixed  solvent  changes. 

Similar  results  were  obtained  with  neodymium  nitrate  and  praseo- 
dymium chloride.  These  results  showed  that  the  solvent  played  an 
important  part  in  the  absorption  of  light  by  solutions.  Jones  and 
Anderson  explained  this  fact  on  the  basis  of  the  solvate  theory  of 
solution  proposed  by  Jones  in  1900,  that  a  part  of  the  solvent  combines 
with  the  dissolved  substance  and  that  the  solvated  parts  have  different 
resonance  from  the  unsolvated. 

The  work  of  Jones  and  Anderson  was  continued  by  Jones  and  Strong,2 
who  studied  a  large  number  of  salts  in  various  solvents,  to  see  if  the 
solvent  in  general  played  a  role  in  the  absorption  of  light  by  solution. 
Most  of  their  work  was  done  with  salts  of  neodymium  and  uranium, 
since  these  substances  had  sharp  absorption  lines  and  bands.  That  the 
solvent  has  much  to  do  with  the  absorbing  power  of  the  solution  can  be 
seen  from  table  53,  which  contains  the  wave-lengths  of  the  absorption 
lines  of  uranyl  chloride  in  various  solvents. 


TABLE  53. 

In  water  

.  .  .  XX4025 

4170 

4315 

4460 

4560 

4740 

4920 

In  methyl  alcohol  

.  .  .  XX4090 

4220 

4345 

4465 

4590 

4760 

4930 

In  ethyl  alcohol  

.  .  .  XX4100 

4250 

4400 

4580 

4750 

4900 

In  propyl  alcohol  

.  .  .  XX4100 

4230 

4400 

4580 

4750 

4910 

In  isopropyl  alcohol 

.  .  XX4100 

4250 

4360 

4560 

4750 

In  butyl  alcohol  

.  .  .  XX4100 

4240 

4390 

4560 

4750 

4970 

In  isobutyl  alcohol  

.  ..  XX.... 

4400 

4560 

4720 

4900 

In  ether  

.  .  .  XX4040 

4160 

4300 

4444 

4630 

In  methyl  ester  

.  .  .  XX4030 

4160 

4280 

4440 

4620 

4990 

4920 

In  glycerol  

.  .  .  XX4025 

4140 

4260 

4400 

4540 

4720 

5050 

In  formamid  

.  ..  XX.... 

4450 

4650 

4840 

'Carnegie  Inst.  Wash.  Pub.  No.  110  (1909);  Amer.  Chem.  Journ,  41,  163  (1909). 

'Carnegie  Inst,  Wash.  Pubs.  Nos.  130  (1910)  and  160  (1911);  Amer.  Chem.  Journ.,  43,  37, 
224  (1910);  45,  1  (1910);  47,27  (1912);  Phys.  Zeit.,  10,  449  (1909);  Phil.  Mag.,  April  1910; 
Journ.  Chim.  Phys.,  8,  131  (1910). 

83 


84  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

Solutions  of  uranyl  chloride  in  such  closely  related  solvents  as  propyl 
and  isopropyl  alcohols  have,  as  is  shown  above,  different  absorption 
lines  and  bands. 

Jones  and  Strong  studied  the  absorption  spectra  of  salts  of  neo- 
dymium,  and  found  different  absorption  lines  and  bands  for  solutions 
of  the  salts  in  isomeric  solvents.  This  is  brought  out  in  table  54. 

TABLE  54. 

In  water XX3390  3465  3505     3540     3560     

In  methyl  and  ethyl  alcohols XX3475  3505  3560     

In  propyl  alcohol XX3545  3460  3490  3510     3525     3540     3560     3580 

In  isopropyl  alcohol XX3460  3510  3535     

In  butyl  alcohol XX3450  3460  3492     3535     3545     3560     

In  isobutyl  alcohol XX3455  3485  3515     3545     3570     

In  glycerol XX3520  3475  3550     

The  nitrate  of  neodymium  was  studied  in  the  same  way  as  the 
chloride.  The  absorption  bands  of  neodymium  nitrate  were  found  to 
be  practically  the  same  as  the  absorption  bands  of  neodymium  chloride. 
This  is  strong  evidence  that  the  solvent  plays  an  important  part  in 
the  absorption  of  light  by  substances  dissolved  in  it. 

Jones  and  Strong  studied  the  effect  of  rise  in  temperature  on  the 
absorption  spectra  of  solutions,  and  found  that  the  absorption  bands 
widened  with  rise  in  temperature.  When  the  solution  is  cooled  down 
the  original  spectrum  is  obtained.  This  was  explained  by  the  initial 
solvates  broken  down  by  rise  in  temperature  being  reformed  on  cooling. 

Jones  and  Guy1  studied  the  effect  of  dilution  on  absorption  spectra, 
and  found  that  the  absorption  bands  widened  as  the  concentration  of 
the  solution  was  increased.  This  is  what  we  should  expect,  since  a 
change  in  the  concentration  of  the  solution  would  mean  a  change  in 
the  complexity  of  the  solvate.  The  more  dilute  the  solution  the  more 
complex  the  individual  aggregate ;  and  this  change  in  complexity  should 
affect  its  power  of  resonance.  Conversely,  the  more  concentrated  the 
solution  the  simpler  each  individual  solvate  and  the  greater  its  power 
of  resonance. 

Jones  and  Guy  studied  the  effect  of  the  dissolved  substance  on  the 
absorption  of  light  by  water.  The  absorptions  of  aqueous  solutions  of 
salts  were  compared  with  the  absorption  of  a  layer  of  water  equal  in 
depth  to  the  water  in  the  solution.  Slightly  hydrated  salts,  such  as 
potassium  chloride  and  ammonium  chloride,  were  found  to  have  about 
the  same  absorption  as  water.  On  the  other  hand,  strongly  hydrated 
salts  like  calcium  and  magnesium  chlorides,  were  found  to  be  far  more 
transparent  than  pure  water.  They  explained  this  fact  on  the  ground 
that  combined  water  has  less  absorption  than  free  water. 

Jones,  Shaeffer,  and  Paulus2  repeated  and  elaborated  the  work  of 
Jones  and  Guy,  and  found  the  above  relations  to  be  general. 

'Carnegie  Inst.  Wash.  Pub.  No.  190  (1913) ;  Ann.  der  Phys.,  43,  555  (1914) ;  Amer.  Chem. 
Journ.,  41,  1  (1913). 

*Phys.  Zeit.,  15,  447  (1914). 


Action  of  Hydrated  and  Nonhydrated  Salts  on  Saponification.  85 

The  above  physical  difference  in  the  action  of  combined  and  of  free 
water,  in  their  power  to  absorb  light,  led  us  to  study  solutions  of 
slightly  hydrated  salts  and  strongly  hydrated  salts  to  see  if  a  corre- 
sponding chemical  difference  existed  between  water  in  the  combined 
and  in  the  free  state — in  other  words,  to  study  the  effect  of  free  water 
on  some  chemical  reaction  in  which  water  is  one  of  the  reacting  sub- 
stances, a  solution  of  a  slightly  hydrated  salt  containing  an  amount  of 
water  equal  to  the  free  water,  and  a  solution  of  a  hydrated  salt  con- 
taining an  amount  of  water  equal  to  the  free  water. 

REACTION  CHOSEN. 

Berthelot  and  Pe"an  de  Saint  Gilles,1  as  early  as  1862,  showed  that  the 
formation  of  an  ester  from  an  alcohol  and  an  acid 

CHsCOOH  +  HOC2H5^CH3COOC2H5+H2O 

proceeded  in  a  slow,  progressive  manner  towards  a  limit,  and  that 
the  speed  of  the  reaction  was  dependent  on  the  amounts  of  the  sub- 
stances present  and  on  the  temperature.  This  was  also  shown  to  be 
a  reversible  process. 

The  reaction  of  the  saponification  of  an  ester  has  been  studied  by 
Ostwald2  and  Reicher,3  as  a  method  for  determining  the  strengths  of 
acids  and  bases.  Bases  were  found  to  saponify  the  ester  more  rapidly 
than  acids.  The  strengths  of  the  acids  and  bases  determined  by  this 
method  agreed  with  the  strengths  as  determined  by  conductivity.  In 
other  words,  it  is  the  hydroxyl  ion  of  the  base  and  the  hydrogen  ion 
of  the  acid  which  do  the  saponifying.  The  velocities  of  the  reactions 
studied  by  Ostwald  and  Reicher  were  so  much  greater,  on  account  of 
the  large  number  of  ions  furnished  by  the  acids  and  bases,  in  contrast 
with  the  small  number  of  ions  in  those  studied  by  Berthelot  and  Saint 
Gilles. 

We  decided  to  use  the  saponification  of  an  ester  in  the  study  of  the 
action  of  free  water  and  of  combined  water  for  the  following  reasons: 

In  the  first  place,  as  Berthelot  and  Saint  Gilles  showed,  the  speed 
of  the  reaction  is  dependent  on  the  relative  amounts  of  the  substances 
used.  By  keeping  the  amount  of  the  ester  constant,  we  could  study 
the  effect  of  the  water  in  solutions  of  salts  containing  the  same  amounts 
of  water. 

Secondly,  the  reaction  proceeds  slowly  and  its  velocity  could  there- 
fore be  easily  measured. 

Thirdly,  the  speed  of  the  reaction  is  dependent  on  the  temperature. 
However,  as  the  temperature  rises  the  hydrates  become  less  complex. 

'Ann.  Chim.  Phys.,  65,  385  (1862);  66,  5,  110  (1862);  68,  225,  (1865);  14  ,437  (1878);  15,  220 
(1878). 
•Jour,  prakt.  Chem.,  28,  449  (1883);  35,  112  (1887).  »Lieb.  Ann.,  228.  257  (1885). 


86  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

From  the  results  obtained  with  the  ester  and  pure  water,  as  compared 
with  those  from  solutions  of  the  salts,  we  could  study  the  effect  of 
rise  in  temperature  on  the  hydrated  salts. 

HISTORICAL. 

The  effect  of  neutral  salts  on  reaction  velocities  was  studied  first 
by  Arrhenius,1  who  investigated  the  hydrolysis  of  ethyl  acetate  as 
effected  by  bases,  and  by  bases  in  the  presence  of  one  of  their  salts. 
He  found  that  in  the  case  of  1  /40  normal  potassium  hydroxide,  sodium 
hydroxide,  and  barium  hydroxide,  the  addition  of  their  halogen  salts 
decreased  the  velocity  of  the  reaction,  and  that  the  addition  of  sulphates 
increased  the  velocity. 

Spohr2  studied  the  effect  of  neutral  salts  on  the  inversion  of  cane 
sugar  by  acids,  and  on  the  saponification  of  an  ester  by  bases.  He 
found  that  the  addition  of  neutral  salts  increased  the  velocity  of  the 
former  reaction,  but  diminished  the  velocity  of  the  latter. 

Arrhenius3  also  studied  the  influence  of  neutral  salts  on  the  inversion 
of*  cane  sugar  by  acids  and  obtained  results  similar  to  those  found  by 
Spohr.  The  effect  of  the  salt  was  greater  with  dilute  than  with  more 
concentrated  acids. 

Euler4  obtained  results  very  similar  to  those  found  by  Spohr. 

The  above  work  led  to  the  study  of  the  effect  of  neutral  salts  when 
the  acid  and  base  were  absent. 

Smith5  measured  the  dissociation  of  organic  dibasic  acids  by  meas- 
uring the  rate  at  which  the  acids  invert  cane  sugar.  He  also  studied 
the  effect  of  neutral  salts  on  the  reaction  when  the  acids  were  absent, 
and  found  that  while  salts  of  weak  acids  did  not  affect  the  reaction 
differently  from  pure  water,  the  salts  of  strong  acids,  such  as  potassium 
chloride  and  sodium  sulphate,  produced  a  large  effect. 

Kellogg6  studied  the  problem  of  the  effect  of  neutral  salts  on  the  re- 
action velocity  of  the  saponification  of  ethyl  acetate.  The  ester  and  the 
solutions  were  sealed  in  glass  tubes  and  put  in  a  constant-temperature 
bath  at  100°.  The  tubes  were  then  opened  and  the  contents  titrated 
with  phenolphthalein  and  sodium  hydroxide.  To  get  the  amounts  of 
the  solutions  of  the  salts  which  should  contain  an  amount  of  water 
equal  to  the  free  water,  the  density  of  the  salt  solution  was  determined, 
its  percentage  of  water  content  figured,  and  the  amount  to  be  added 
calculated.  Kellogg  found  that  in  the  more  dilute  solutions  the  specific 
influence  of  the  salt  is  greater.  As  the  solutions  become  more  concen- 
trated the  increase  in  reaction  velocity  grows  less,  and  in  the  case  of 
4-normal  potassium  chloride  the  hydrolysis  is  slower  than  in  pure 
water.  He  found  a  decrease  in  the  accelerating  power  from  potassium 
chloride,  to  potassium  bromide,  to  potassium  iodide,  which  is  in  the 

'Zeit.  phys.  Chem.,  1,  110  (1887).         '/few*.,  4,  237  (1889).         '/bid.,25, 144 (1898). 
'Ibid.,  2,  194  (1888).  *Ibid.,  32, 348  (1900).         Mourn.  Amer.  Chem.  Soc.,  31, 

403,  886  (1909). 


Action  of  Hydrated  and  Nonhydrated  Salts  on  Saponification.  87 

reverse  order  of  their  stability.  The  curves  plotted  with  time  and 
percentage  hydrolysis  as  axes  are  of  the  same  general  character,  with  a 
concentration  of  maximum  accelerating  power,  which  is  1 .8  normal  for 
potassium  chloride,  0.5  normal  for  potassium  bromide,  and  0.25  normal 
for  potassium  iodide.  Abnormally  low  results  were  found  with  potas- 
sium iodide,  which  he  explained  as  due  to  the  decomposition  of  the 
potassium  iodate  present  by  the  acetic  acid  formed  in  the  reaction 

KIO3+  6KI  +  6CH3COOH  -  -  KI  +3I2+6CH3COOK  +3H20 

These  solutions  were  found  to  be  slightly  colored,  due  to  the  iodine 
set  free.  By  varying  the  time,  Kellogg  studied  the  reaction  from  the 
point  where  1  per  cent  of  the  ester  was  hydrolyzed  to  practically  com- 
plete hydrolysis.  The  salt  has  two  effects,  the  one  in  dilute  solution 
being  accelerating,  the  other  increasing  as  the  concentration  increases, 
retarding  the  reaction. 

Kellogg  made  three  suggestions  to  explain  this  phenomenon.  One 
hypothesis  is  that  the  accelerating  effect  is  due  to  the  ionized  salt 
and  the  retardation  is  due  to  the  undissociated  salt.  This  would 
explain  the  decreasing  acceleration  with  increasing  concentration,  as 
the  ionized  part  increases  slowly  in  proportion  to  the  non-ionized  part. 
It  would  also  account  for  the  facts  found  by  the  earlier  investigators, 
that  the  salt  of  an  acid  has  the  greatest  effect  with  dilute  acids,  since 
with  a  dilute  acid  we  have  a  smaller  number  of  common  ions  and 
therefore  a  greater  dissociation  of  the  salt. 

Kellogg  also  stated  that  the  two  effects  might  be  functions  of  the 
ions;  one  of  the  potassium,  the  other  of  the  halogen  ion.  This  would 
account  for  the  difference  between  the  three  salts  which  he  used,  but 
not  for  the  difference  between  weak  and  strong  solutions. 

Euler's1  suggestion  would  explain  the  dual  effect.  He  assumes  that 
the  salt  increases  the  reactive  capacity  of  the  water;  the  ions  of  the 
neutral  salt  combining  with  the  solvent.  The  water  around  the  ion 

of  the  salt  is  at  a  high  tension,  and  we  therefore  would  have  a  large 

+ 

number  of  H  and  OH  ions.  This  would  explain  the  larger  effect  of 
neutral  salts  with  dilute  acids.  Spohr  and  Arrhenius  held  practically 
the  same  views  as  Euler. 

Euler  also  suggested  that  intermediate  compounds  might  be  formed. 

Arrhenius's  older  idea  of  the  "active"  and  "inactive"  forms  of  the 
ester  would  explain  all  of  the  facts,  the  salt  disturbing  the  equilibrium 
in  favor  of  the  "active"  form. 

In  conclusion,  Kellogg  stated  that  no  one  theory  could  explain  all 
of  the  facts,  each  having  its  merits  and  its  disadvantages. 

Henderson  and  Kellogg,2  continuing  the  work  of  Kellogg,  studied  the 
chlorides  of  sodium,  lithium,  calcium,  strontium,  and  barium,  and  the 

'Zeit.  phys.  Chem.,  32,  348  (1900).  'Journ.  Amer.  Chem.  Soc.,  35,  396  (1913). 


88  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

chloride  and  iodide  of  cadmium.  They  measured  the  conductivity 
and  viscosity  of  the  solutions  at  the  concentrations  and  temperature 
employed  in  the  experiments,  and  from  these  calculated  the  degree  of 
ionization. 

The  work  was  carried  out  under  the  same  conditions  employed  by 
Kellogg.  They  found  that  all  of  the  salts  used,  even  at  concentrations 
as  low  as  0.1  normal,  increase  the  rate  of  the  reaction.  As  the  concen- 
tration of  the  salt  increases  the  effect  becomes  less  and,  finally,  when 
the  solution  becomes  concentrated,  the  salt  has  a  retarding  effect. 

The  curves  they  obtained  for  these  salts  were  found  to  be  very  similar 
to  those  for  the  alkali  chlorides,  but  showed  a  larger  effect.  Cadmium 
chloride  and  iodide  were  found  to  give  a  very  large  value,  due  probably 
to  their  hydrolysis,  since  they  have  an  acid  reaction. 

Henderson  and  Kellogg  found  that  the  salts  which  produced  the 
greatest  effect  were  those  that  were  least  ionized.  From  this  they 
drew  the  conclusion  that  the  solution  of  the  problem  was  to  be  sought 
in  the  effect  of  the  non-ionized  portion  of  the  salt  rather  than  in  the 
ionized  portion.  If  the  hydrolysis  of  the  ester  is  controlled  by  the 
concentration  of  the  salt  molecules,  the  cause  of  the  observed  maximum1 
is  easily  explained,  since  as  the  total  concentration  of  the  salt  increases, 
not  only  does  the  concentration  of  the  molecules  increase,  but  the  vis- 
cosity also  varies.  The  maximum,  then,  was  the  resultant  between  the 
acceleration  produced  by  the  salt  molecules  which  increases  with  the 
concentration,  and  the  resistance  to  the  contact  of  the  ester  and  water, 

due  to  the  increasing  viscosity. 

+ 

Measurements  of  the  H  and  OH  ions,  had  they  been  carried  out, 
would  have  had  an  important  bearing  on  Euler's2  hypothesis;  that  at 
higher  salt  concentrations  water  is  much  more  ionized  than  in  dilute 
solutions,  and  it  is  the  increase  in  the  ions  from  water  that  accelerates 
the  reaction. 

HYDROLYSIS. 

The  presence  of  hydrogen  or  hydroxyl  ions  other  than  those  from 
water  must  be  guarded  against  with  greatest  care.  Since  the  ions 
from  water  are  relatively  so  few  in  number,  the  introduction  of  ions 
from  other  sources  would  render  the  results  of  little  value.  Ions  might 
easily  be  introduced  in  the  ester  and  in  the  solutions  of  the  salts.  The 
ester  must  be  distilled  until  it  gives  a  neutral  reaction.  Salts  of  weak 
acids  or  weak  bases  can  not  be  used  on  account  of  hydrolysis: 

KCN  T±  K  +  CN 
HOH  ^0+H  +  H 


KOH  HCN 


"Journ.  Amer.  Chem.  Soc.,  31,  886  (1909).  JZeit.  phys.  Chem.,  32,  348  (1900). 


Action  of  Hydrated  and  Nonhydrated  Salts  on  Saponification.  89 

The  hydrocyanic  acid  would  dissociate  very  slightly,  while  the  potas- 
sium hydroxide  would  be  almost  entirely  dissociated,  giving  a  large 

excess  of  OH  ions,  and  therefore  an  alkaline  reaction.  The  hydroxyl 
ions,  being  in  large  numbers  as  compared  with  the  hydroxyl  ions  from 
the  water,  would  vitiate  the  results.  In  the  case  of  a  salt  of  a  weak 
base  we  would  find  an  acid  reaction. 

The  hydrolysis  of  salts  can  be  calculated  from  the  formula  of 
Arrhenius,1 

(Salt) Kacid 

(Acid)  X  (Base)  ==  HKOH 

in  which  the  salt,  acid,  and  base  represent  the  total  concentrations, 
ionized,  and  non-ionized,  at  the  point  of  equilibrium;  and  the  K's  are 
the  ionization  constants.  Similar  formulae  were  worked  out  by  Walker2 
for  a  salt  of  a  weak  base  and  a  strong  acid,  and  by  Arrhenius  for  a  salt 
of  a  weak  acid  and  a  weak  base. 

Where  the  ionization  constant  can  not  be  determined  by  conduc- 
tivity measurements,  other  methods  must  be  used. 

Shields3  employed  the  saponification  of  ethyl  acetate  and  showed 
that  with  increasing  dilution  the  hydrolysis  increased. 

Ley4  used  the  saponification  of  methyl  acetate,  the  inversion  of 
cane  sugar,  and  the  conductivity  methods;  he  found  that  these  three 
methods  gave  results  which  agreed  very  well  with  one  another.  These 
being  the  most  general  methods,  are  the  only  ones  that  need  to  be 
considered. 

Since  the  reaction  of  the  saponification  of  an  ester  has  been  used  to 
measure  the  hydrolysis  of  salts  such  as  potassium  cyanide  and  aluminum 
chloride,  which  react  basic  and  acidic  with  indicators,  it  is  of  primary 
importance  that  neutral  salts  be  used. 

Salts  which  give  no  reaction  with  litmus,  according  to  Salm,6  have 

+ 

a  concentration  of  H  or  OH  ions  less  than  0.000001,  which  is  so  small 
that  it  is  practically  negligible. 

We  found  that  the  salts  of  potassium,  sodium,  lithium,  magnesium, 
calcium,  barium,  and  strontium,  with  chlorine,  bromine,  iodine  and  the 
nitric  and  sulphuric  acid  anions,  gave  no  reaction  with  litmus.  The 
acids  have  an  ionization  constant  which  is  very  nearly  unity ;  and  the 
weakest  of  the  bases  has  so  large  a  constant  that  if  substituted  in 
Walker's  formula  the  hydrolysis  would  be  negligible. 

Finally,  Bruner6  showed  that  the  hydrolysis  of  solutions  of  the 
chlorides  of  potassium,  lithium,  barium,  strontium,  calcium,  and  mag- 
nesium was  too  small  to  be  measured.  He  also  pointed  out  that,  in  general, 
chlorides  were  most  hydrolyzed,  then  nitrates,  and  least  of  all  sulphates. 

Therefore,  in  the  solutions  to  be  studied  the  effect  due  to  hydrolysis 
may  be  neglected. 

'Zeit.  phys.  Chem.,  5,  16  (1890).    'Ibid.,  12,  167  (1893).     'Ibid., 57, 471  (1907). 
'Ibid.,  4,  319  (1889).  4/Wd.,  30,  193  (1899).     '/bid.,  32,  133  (1900). 


90  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

STATEMENT  OF  PROBLEM. 

The  plan  was  to  investigate  the  difference  in  the  velocities  with 
which  free  water  and  combined  water  saponified  an  ester  under  the 
following  conditions : 

(1)  Time  and  concentration  of  the  salts  kept  constant,  temperature 
varying. 

(2)  Concentration   of  the  salts  and  temperature  constant,   time 
varying. 

(3)  Time  and  temperature  constant,  concentration  of  salts  varying. 

The  conditions  were  so  regulated  that  the  percentage  of  ester  saponi- 

+ 

fied  was  always  small,  the  H  ions  of  the  acid  formed,  being  in  rela- 

+ 

tively  large  numbers  as  compared  with  the  H  ions  from  water,  would 
suppress  the  ionization  of  water  and  themselves  effect  the  saponification. 

EXPERIMENTAL. 
APPARATUS. 

The  reactions  were  allowed  to  take  place  in  100  c.c.  Jena  bottles 
with  ground-glass  stoppers.  These  bottles  were  suspended  in  constant- 
temperature  baths,  very  similar  to  those  employed  by  Davis  and 
Putnam.1  The  thermometers  were  compared  with  a  standardized 
thermometer.  The  temperatures  of  the  baths  were  kept  constant  to 
within  0.02  of  a  degree.  The  baths  were  stirred  at  first  by  a  gas- 
engine,  later  by  an  electric  motor.  The  solutions  were  measured  by 
means  of  carefully  calibrated  pipettes;  10  c.c.  burettes  were  used  for 
measuring  the  ester,  and  50  c.c.  burettes  for  titration  purposes.  All 
measuring  flasks  were  made  of  Jena  glass  and  were  calibrated  by  weight. 
The  solutions  were  always  kept  in  Jena  bottles. 

SOLUTIONS. 

As  stated  before,  only  solutions  which  gave  neutral  reactions  with 
litmus  were  us"ed.  The  salts  were  Kahlbaum  preparations  recrystal- 
lized.  The  solutions  of  slightly  hydrated  salts  were  made  up  by  weight, 
the  strongly  hydrated  salts  by  analysis  and  dilution  to  the  required 
normality. 

A  concentration  was  chosen  arbitrarily,  such  that  solutions  of  all  the 
salts  to  be  investigated  could  be  obtained,  i.  e.,  2  normal.  Later  it  was 
found  to  be  necessary  to  change  this  concentration  to  normal,  using 
also  the  more  dilute  solutions,  half-normal  and  quarter-normal.  The 
amount  of  the  solution  that  contained  an  amount  of  water  equal  to  the 
free  water  was  determined  as  follows:  10  c.c.  of  the  solution  was  weighed 
and  from  this  the  weight  of  1  c.c.  was  determined.  The  weight  of  the 

'Carnegie  Inst.  Wash.  Pub.  No.  210,  117  (1915). 


Action  of  Hydrated  and  Nonhydrated  Salts  on  Saponification.  91 

salt  per  cubic  centimeter  was  subtracted  from  this  weight,  giving  the 
weight  of  the  water  per  cubic  centimeter  of  the  salt  solution.  By 
taking  the  reciprocal  of  this,  the  amount  of  the  solution  containing 
an  amount  of  water  equal  to  1  c.c.  was  obtained.  The  same  trouble 
was  found  with  solutions  of  potassium  iodide  and  sodium  iodide  being 
colored,  as  Kellogg1  describes.  A  drop  of  sodium  thiosulphate  was 
used  to  decolorize  the  solution  before  titrating,  as  the  color  obscured 
the  change  in  the  indicator.  All  solutions  of  the  halogens  were  stand- 
ardized as  their  silver  salts.  Calcium  nitrate  was  converted  into  the 
oxide,  magnesium  nitrate  into  the  pyrophosphate,  strontium  nitrate 
into  the  sulphate,  and  lithium  nitrate  into  the  sulphate  by  evaporating 
with  sulphuric  acid.  Magnesium  sulphate  was  standardized  as  barium 
sulphate,  and  the  iodides  of  sodium  and  potassium  were  also  determined 
gravimetrically. 

THE  ESTERS. 

The  ester  first  employed  was  ethyl  acetate.  After  some  preliminary 
work,  this  was  discarded  in  favor  of  methyl  acetate,  since  its  solubility 
in  some  of  the  strong  salt  solutions  was  so  slight  as  to  give  us  a  hetero- 
geneous system.  Methyl  formate  was  also  studied,  to  see  if  the  results 
obtained  with  methyl  acetate  were  of  a  general  character  or  specific 
to  the  ester  in  question.  The  esters,  which  were  Kahlbaum  prepara- 
tions, were  distilled  several  times  until  they  boiled  at  the  proper  tem- 
perature and  their  reaction  was  neutral.  The  methyl  formate  had  to 
be  distilled  frequently,  on  account  of  its  apparent  instability. 

THE  BASE  AND  THE  INDICATOR. 

The  finding  of  the  proper  base  and  indicator  to  be  used  in  titrating 
the  acid  formed  gave  us  quite  a  little  trouble.  Standard  sulphuric 
acid  was  made  up,  and  on  analysis  found  to  contain  0.01205  gram  of 
sulphuric  acid  per  cubic  centimeter.  This  was  used  to  standardize  the 
base,  the  comparison  being  repeated  from  time  to  time  to  see  that  the 
base  did  not  undergo  any  change  in  concentration.  The  acid  was  also 
restandardized  from  time  to  time,  no  change  in  its  concentration  being 
noticed.  The  alkali  used  first  was  alcoholic  sodium  hydroxide;  the 
indicator  was  phenolphthalein.  The  precipitate  of  sodium  carbonate 
was  filtered  from  the  alcoholic  solution  in  an  atmosphere  free  from 
carbon  dioxide,  by  a  method  similar  to  that  employed  by  Morse.2  The 
alkali  was  one-tenth  normal.  Great  difficulty  was  experienced  in 
obtaining  good  end-points  with  the  above  base  and  indicator.  This 
might  be  due  to  the  saponification  of  the  ester  by  the  strong  base,  or  to 
the  unfitness  of  the  indicator.  We  therefore  decided  to  test  out  dif- 
ferent indicators  and  bases. 

'Journ.  Amer.  Chem.  Soc.,  31,  886  (1909).  'Exercises  in  Quant.  Chem.,  330. 


92  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

We  obtained  the  following  results: 

Ammonia  and  phenolphthalein: 

10  e.c.  standard  acid  in  presence  of  the  ester — -19.3  c.c.  of  base  used. 

10  c.c.  standard  acid — 16.3  c.c.  of  base. 
Sodium  hydroxide  and  phenolphthalein: 

5  c.c.  HjSOi,  strength  unknown  in  presence  of  the  ester — 16.8  c.c.  of  base  used. 

5  c.c.  H2SO4,  strength  unknown — 13.60  c.c.  of  base. 
Ammonia,  and  methyl  orange: 

5  c.c.  standard  acid  in  presence  of  the  ester — 7.8  c.c.  of  base  used. 

5  c.c.  standard  acid — 7.82  c.c.  of  base. 
Sodium  hydroxide  and  methyl  orange: 

5  c.c.  standard  acid  in  presence  of  the  ester — 12.60  c.c.  of  base  used. 

5  c.c.  standard  acid — 12.55  c.c.  of  base. 
Sodium  hydroxide  and  methyl  orange: 

5  c.c.  HijSOi,  strength  unknown  in  presence  of  the  ester — 14.5  c.c.  of  base  used. 

5  c.c.  H2SO<,  strength  unknown — 13.5  c.c.  of  base. 
Ammonia  and  methyl  orange: 

5  c.c.  H2SO4,  strength  unknown  in  presence  of  the  ester — 8.7  c.c.  of  base  used. 

5  c.c.  HjSOi,  strength  unknown — 8.6  c.c.  of  base. 
Ammonia  and  phenolphthalein: 

5  c.c.  H2SO4,  strength  unknown  in  presence  of  the  ester — 9.4  c.c.  of  base  used. 

5  c.c.  H2SO4,  strength  unknown — 9.1  c.c.  of  base. 

Thus  far,  ammonia  and  methyl  orange  gave  the  best  results.  With 
sodium  hydroxide  the  color  of  the  indictaor  rapidly  disappears  on 
standing,  due  probably  to  saponification  of  the  e«ter.  Barium  hydrox- 
ide has  frequently  been  used  for  this  purpose;  therefore  we  tried  this 
base. 

Barium  hydroxide  and  phenolphthalein: 

5  c.c.  HjSOi,  strength  unknown  in  presence  of  the  ester — 6.0  c.c.  of  base  used. 

5  c.c.  HjSC^,  strength  unknown — 5.8  c.c.  of  base. 
Barium  hydroxide  and  methyl  orange: 

5  c.c.  HjSO*,  strength  unknown  in  presence  of  the  ester — 5.8  c.c.  of  base  used. 

5  c.c.  HjSOi,  strength  unknown — 5.7  c.c.  of  base. 

The  color  of  the  indicator  did  not  persist  with  barium  hydroxide. 
Therefore,  we  discarded  this  in  favor  of  ammonium  hydroxide.  As 
regards  acetic  acid,  methyl  orange  is  the  only  common  indicator  that 
can  be  used. 

Ammonia  and  methyl  orange: 

10  c.c.  acetic  acid,  strength  unknown  in  presence  of  ester — 0.2  c.c.  of  base  used. 
10  c.c.  acetic  acid,  strength  unknown — 0.1  c.c.  of  base. 

Methyl  orange  is  of  little  value,  since  we  should  have  obtained 
readings  of  about  8  or  10  c.c.  Morse  suggests  carollin,  since  this  is  a 
good  indicator  with  organic  acids.1 

Ammonia  and  corattin: 

10  c.c.  acetic  acid,  strength  unknown  in  presence  of  ester — 
9.0  c.c.  base  used. 

9.0  c.c.  base  used. 

10  c.c.  acetic  acid,  strength  unknown — 

9.1  c.c.  of  base. 
9.1  c.c.  of  base. 

'Morse:  Exercises  in  Quant.  Chem.,  289;  Glaser:  Indikatoren,  92. 


Action  of  Hydrated  and  Nonhydrated  Salts  on  Saponification.  93 

To  see  if  the  addition  of  salts  would  have  any  effect  on  the  titration, 
we  titrated  in  the  presence  of  the  ester  and  a  salt : 

10  c.c.  acetic  acid,  strength  unknown: 

In  presence  of  the  ester  and  calcium  chloride;  9.05  c.c.  of  base  used. 
In  presence  of  the  ester  and  potassium  chloride;  9.0  c.c.  of  base  used. 

Ammonium  hydroxide  and  alcoholic  corallin  are,  then,  the  base  and 
the  indicator  to  be  used. 

At  first  one-fifth  normal  ammonia  was  used,  but  this  gave  us  such 
small  volumes  that  we  decided  to  use  one-twentieth  normal. 

After  preparing  the  solution  of  the  base,  it  was  found  to  be  not  quite 
one-twentieth  normal,  but  to  contain  0.00167732  gram  per  c.c. 

For  use  with  the  methyl  formate  a  fifth-normal  solution  was  made  up 
which  contained  0.007  gram  per  c.c. 

To  see  if  the  methods  we  were  using  were  accurate,  the  following 
experiments  were  carried  out : 

0.0733  gram  of  glacial  acetic  acid  was  weighed  out  and  titrated  with 
ammonia,  corallin  being  used  as  the  indicator.  The  amount  of  acetic 
acid  calculated  from  the  titration  was  0.0730  gram.  We  would  then 
have  an  error  of  less  than  0.5  per  cent.  With  other  organic  acids  we 
found  errors  of  less  than  0.1  per  cent. 

The  results  obtained  with  formic  acid  of  about  98.5  per  cent  purity 
were  0.4917  gram  weighed  out,  and  0.4842  gram  was  calculated  from  the 
titration,  giving  us  an  error  of  about  0.1  per  cent. 

METHOD  OF  PROCEDURE. 

Solutions  of  the  salts  and  the  ester  were  measured  into  bottles, 
which  were  then  placed  in  the  baths.  After  the  required  time  had 
elapsed,  the  bottles  were  removed  from  the  baths,  and  their  contents 
titrated  against  the  standard  base.  We  found,  after  a  good  deal  of 
preliminary  experimenting,  that  the  most  favorable  conditions  for 
work  were:  for  the  methyl  acetate,  4  c.c.  of  the  ester  to  30  c.c.  of  water, 
24  and  48  hours,  15,  25,  and  35  degrees,  and  normal,  half -normal,  and 
quarter-normal  salt  solutions.  For  the  methyl  formate,  2  c.c.  of  the 
ester  to  30  c.c.  of  water,  4  and  8  hours,  15  and  25  degrees  (b.  p.  of 
methyl  formate — 32.3°)  and  the  same  concentrations  of  the  salts  as 
were  used  with  the  methyl  acetate.  The  bottles,  when  taken  out  of 
the  baths,  were  immediately  chilled  to  retard  the  reaction.  The  time 
required  for  titrating  was  practically  the  same  as  that  for  adding 
the  ester;  therefore,  the  results  are  comparable.  12  solutions  were 
studied  at  a  time,  6  being  duplicates;  2  bottles  containing  pure  water 
and  ester  were  included  in  the  sets  of  12  bottles,  to  act  as  standards 
and  to  show  that  the  conditions  did  not  vary.  The  addition  of  the 
ester  and  the  titration  of  the  contents  of  the  bottles  were  always 
carried  on  in  the  same  order;  e.  g.,  the  bottle  to  which  the  ester  was 
added  was  titrated  first.  The  methyl  formate  had  to  be  measured  at 


94  Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

15°,  on  account  of  its  high  vapor-tension.  At  first  the  titration  was 
made  in  beakers;  later  it  was  found  that  the  titration  could  be  carried 
on  very  well  in  the  bottle  itself,  thus  eliminating  a  source  of  error.  An 
electric  motor  was  used  for  stirring  the  baths,  since  the  hot-air  engine 
could  not  be  run  over  night.  When  the  baths  were  not  stirred  during 
the  night,  they  varied  as  much  as  0.3°  or  0.4°;  and  the  position 
of  the  bottles  as  regards  the  heating  and  cooling  surfaces  seemed  to 
produce  a  marked  effect.  Normal  solutions  of  sodium  and  potassium 
sulphates  and  of  barium  nitrate  could  not  be  prepared,  since  these 
salts  were  not  sufficiently  soluble;  therefore,  they  were  not  studied. 
Cadmium  and  aluminum  salts  were  discarded  on  account  of  their  acid 
reaction,  and  ammonium  salts  because  they  suppressed  to  nearly  half- 
value  the  ionization  of  the  alkali  used  in  the  titration.  The  ammonium 
salt  formed  by  the  neutralization  of  the  acid  titrated  with  the  base 
would  also  affect  the  ionization  of  the  base,  if  the  reaction  was  allowed 
to  proceed  very  far. 

CALCULATIONS. 

The  results  are  expressed  in  percentages  of  the  ester  decomposed. 
This  was  obtained  as  follows: 

In  the  first  place  the  total  amount  of  the  ester  per  bottle  was  deter- 
mined. From  the  mean  weight  of  four  10  c.c.  portions  the  weight  per 
cubic  centimeter  of  the  methyl  acetate  was  found  to  be  0.9257  gram, 
giving  us  3.7028  grams  as  the  total  weight  of  the  methyl  acetate. 

Then  the  weight  of  the  ester,  which,  on  saponification,  would  give 
enough  acetic  acid  to  neutralize  1  c.c.  of  the  standard  one-twentieth 
normal  ammonia,  was  calculated;  1  c.c.  of  ammonia  corresponded  to  a 
decomposition  of  0.003543  gram  of  methyl  acetate.  Dividing  this 
value  by  the  total  amount  of  the  ester,  3.7028  grams,  we  find  that  1  c.c. 
of  one-twentieth  normal  ammonia  (0.001677  gram  per  c.c.)  represents 
0.0957  per  cent  of  the  ester  saponified. 

To  get  the  percentage  saponification  for  each  salt,  it  is  only  necessary  to 
multiply  the  burette  reading  for  each  bottle  by  the  factor  0.0957  per  cent. 

For  methyl  formate  we  obtained  the  following  values:  The  weight 
per  c.c.  of  ester,  0.9701  gram;  the  total  weight  of  the  ester,  1.9402 
grams;  the  weight  of  the  ester  which,  when  saponified,  corresponded 
to  1  c.c.  of  the  fifth-normal  ammonia  (per  c.c.  0.007  gram)  is  0.01199 
gram,  and  the  percentage  saponification,  0.618  per  cent. 

DATA. 

In  the  following  tables  of  data,  the  methyl  acetate  per  bottle  is  4  c.c. 
or  3.7028  grams;  the  methyl  formate,  2  c.c.  or  1.9402  grams.  Table 
56  contains  the  results  when  the  baths  were  not  stirred  over  night,  and 
these  results  have,  therefore,  a  slight  error.  It  was  found  that  calcium 
chloride  and  potassium  nitrate,  on  dilution,  increase,  then  decrease  the 
velocity  of  the  reaction.  Duplicate  measurements  were,  therefore, 
made  with  new  solutions,  and  the  results  recorded  in  table  58.  The 


Action  of  Hydrated  and  Nonhydrated  Salts  on  Saponifization.  95 


data  in  tables  64  and  65  were  obtained,  using  the  same  salt  solutions. 
It  was  at  this  point  that  we  decided  to  study  also  methyl  formate. 
For  tables  67  and  69  new  solutions  were  used,  since  those  earlier  em- 
ployed had  not  been  preserved.  The  volumes  of  the  solutions  of  the 
salts  that  contained  30  c.c.  of  water  were  calculated  as  described 
above,  the  water  per  cubic  centimeter  of  the  solution  being  the  differ- 
ence in  weight  between  1  c.c.  of  the  solution  and  the  salt  in  1  c.c.  The 
water  used  in  preparing  the  solutions  of  the  salts,  bases,  and  acid  had 
been  carefully  purified,  and  had  a  mean  specific  conductivity  of  1.5  X  10~6 
at  25° .  This  water  was  also  employed  hi  recry stallizing  the  salts.  In  the 
tables  of  data  the  concentrations  of  the  solutions  of  the  salts  are  ex- 
pressed as  N,  N/2,  and  N/4.  Tables  56,  58,  60,  62,  and  64  give  the 
percentages  of  the  methyl  acetate  saponified;  tables  65,  67,  and  69  of 
the  methyl  formate.  Duplicate  measurements  were  made  with  fresh 
solutions  of  magnesium  sulphate  and  strontium  bromide;  since  the 
results  obtained  with  these  salts  were  very  low  as  compared  with  the 
results  from  the  other  salts  that  contain  a  large  amount  of  water  of 
crystallization.  The  results  found  were,  however,  identical  with  those 

earlier  obtained. 

TABLE  55. 


Volume 

Volume 

Volume 

Volume 

Grams  of 

of  solu- 

of solu- 

Grams of 

of  solu- 

of solu- 

Solutions of  the 

H2O  per 

tion 

tion 

Solutions  of  the 

H2O  per 

tion 

tion 

salts. 

e.c.  of  the 

cont.1 

cont. 

salts. 

c.c.  of  the 

cont. 

cont. 

solution. 

1  c.c. 

30  c.c. 

solution. 

1  c.c. 

30  c.c. 

H2O. 

H20. 

H2O. 

H2O. 

N      KC1.    .. 

0  9717 

1  029 

30  87 

N/2  MgCl2   

0.9874 

1.013 

30.39 

N      KNO, 

0  9610 

1  041 

31  23 

N  /2  SrCla 

0  9852 

1  015 

30  45 

N      CaCli 

0  9721 

1  028 

30  84 

N/4  KC1  

0.9910 

1.009 

30.27 

N      MgClj 

0  9806 

1  020 

30  60 

N/4  KNOa      

0.9845 

1.016 

30.48 

N      SrCl2  

0  9716 

1  029 

30  87 

N/4  CaClj  

0.9921 

1.008 

30.24 

N  /2  KC1 

0  9846 

1  016 

30  48 

N/4  MgClj     

0.9924 

1.008 

30.24 

N/2  KNOa 

0  9789 

1  022 

30  66 

N  /4  SrCh 

0  9909 

1.009 

30.27 

N/2  CaCl2  

0.9871 

1.013 

30.39 

TABLE  56. 


No. 

Solutions 

N. 

15° 
24  hrs. 

25° 
24  hrs. 

35° 
24  hrs. 

25° 
48  hrs. 

Solutions 

N/2. 

28° 

24  hrs. 

35° 

24  hrs. 

Solutions 

N/4. 

25° 

24  hrs. 

1 
2 

f  30  c.c.  1 
I   H20.     ] 

0.076 
0.076 

0.147 
0.145 

0.390 
0.392 

0.287 
0.289 

/  30  c.c.    1 
I  H20.     j 

0.145 
0.144 

0.386 
0.385 

f  30  c.c.   \ 
I  H20.     / 

0.153 
0.152 

3 

/30.87  o.c.\ 

0.100 

0.224 

0.720 

0.516 

/30.48  c.c.1 

0.206 

0.669 

/30.27  c.c.1 

0.196 

4 

1   KC1.     ] 

0.100 

0.222 

0.722 

0.513 

1    KC1.      j 

0.206 

0.666 

1   KC1.     / 

0.198 

5 

/31.23  e.c.1 

0.099 

0.194 

0.611 

0.431 

/30.66  c.c.1 

0.201 

0.612 

/30.48  c.c.1 

0.177 

6 

IKNO,.  / 

0.097 

0.196 

0.614 

0.431 

IKNO,.  ] 

0.201 

0.613 

IKNO,.   / 

0.177 

7 

/30.84  c.c.\ 

0.119 

0.387 

1.68 

1.03 

/30.39  c.c.1 

0.400 

1.35 

/30.24  c.c.1 

0.278 

8 

1  CaCl,.    j 

0.117 

0.385 

1.69 

1.05 

\CaClj.    J 

0.398 

1.36 

\CaCl2.    J 

0.277 

9 

/30.60  c.c.\ 

0.132 

0.410 

1.67 

1.33 

/30.39  c.c.1 

0.326 

1.23 

/30.24  c.c.\ 

0.276 

10 

\MgClj.   / 

0.134 

0.411 

1.67 

1.30 

\MgCl,.  / 

0.323 

1.23 

\MgCl,.  } 

0.276 

11 

/30.87  c.c.1 

0.130 

0.378 

1.41 

0.992 

/30.45  c.c.1 

0.297 

1.12 

f  30.27  c.c.\ 

0.236 

12 

\  SrCl,.    ] 

0.128 

0.375 

1.42 

0.989 

I  SrCU.    ] 

0.297 

1.12 

I  SrClj.    / 

0.239 

1  "Cont."  in  this  and  the  following  tables  is  "containing." 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


TABLE  57. 


Solutions  of  the 

salts. 

Grams  of 
H2O  per 
c.c.  of  the 
solution. 

Volume 
of  solu- 
tion 
cont. 
1  c.c. 
H,O. 

Volume 
of  solu- 
tion 
cont. 
30  c.c. 
H20. 

Solutions  of  the 
salts. 

Grams  of 
H2O  per 
c.c.  of  the 
solution. 

Volume  Volume 
of  solu-'  of  solu- 
tion   j    tion 
cont.    '   cont. 
1  c.c.   I  30  c.c. 
H2O.       H,O. 

N      LiCl 

0.9688 
0.9589 
0.9811 
0.9718 
0.9496 
0.9840 
0.9802 
0.9903 

1.032 
1.043 
1.019 
1.029 
1.053 
1.016 
1.020 
1.010 

30.96 
31.29 
30.57 
30.87 
31.59 
30.48 
30.60 
30.30 

N  /2  CaCh 

0.9868 
0.9773 
0.9917 
0.9870 
0.9931 
0.9909 
0.9880 

1.013 
1.023 
1.008 
1.013 
1.007 
1.009 
1.012 

30.39 
30.69 
30.24 
30.39 
30.21 
30.27 
30.36 

N      KNO>   . 

N/2Ca(NO3)2  
N  /4  LiCl 

N      NaCl  .    ... 

N      CaCl2  

N  /4  KNOj  .  . 

N      Ca(NOs)j  
N  /2  LiCl 

N/4NaCl  

N/4  CaCU 

N/2  KNOs  

N/4  Ca(NO3)2 

N/2  NaCl  

TABLE  58. 


No. 

Solutions 

N. 

15° 

24  hrs. 

25° 
24  hrs. 

35° 
24  hrs. 

25° 
48  hrs. 

Solutions 

N/2. 

25° 
24  hrs. 

35° 
24  hrg. 

Solutions 

N/4. 

25° 
24  hrs. 

1 
2 

/  30  c.c.    "1 
1   H,0.     ] 

0.083 
0.084 

0.154 
0.154 

0.394 
0.392 

0.296 
0.297 

1  30  c.c.    1 
1   H80.     / 

0.153 
0.153 

0.391 
0.390 

/  30  c.c.    \ 
I   H20.     I 

0.152 
0.151 

3 
4 

[30.96  c.c.1 
{   LiCl.     / 

0.094 
0.093 

0.204 
0.203 

0.626 
0.624 

0.462 
0.464 

/30.48  c.c.\ 
1   LiCl.     / 

0.199 
0.198 

0.590 
0.591 

/30.24  c.c.l 
1    LiCl.     / 

0.194 

0.194 

5 
6 

/31.29  c.c.\ 

IKNO,.  / 

0.099 
0.100 

0.215 
0.214 

0.631 
0.633 

0.457 
0.456 

/30.60  c.c.\ 

IKNO,.  / 

0.213 
0.212 

0.632 
0.634 

/30.39  c.c.l 
\KN03.   / 

0.212 
0.211 

7 
8 

/30.57  c.c.1 
\  NaCl.    / 

0.109 
0.111 

0.247 
0.245 

0.808 
0.806 

0.590 
0.591 

/30.30  c.c.\ 
1  NaCl.    / 

0.241 
0.242 

0.746 
0.745 

/30.21  c.c.l 
1  NaCl.     / 

0.222 
0.222 

9 
10 

f  30.87  c.c.\ 
ICaCl,.    / 

0.129 
0.131 

0.393 
0.395 

1.72 
1.71 

1.15 
1.16 

/30.39  c.c.l 
\  CaCIj.    / 

0.306 
0.308 

1.35 
1.36 

/30.27  c.c.l 
1  CaC2.     / 

0.285 
0.284 

11 
12 

/31.59  c.c.l 
lCa(NO,)«j 

0.129 
0.126 

0.330 
0.333 

1.32 
1.34 

0.860 
0.861 

/30.69  c.c.l 
lCa(NO,)2/ 

0.310 
0.308 

1.12 
1.12 

/30.36  c.c.1 

lCa(N03)2/ 

0.270 
0.269 

TABLE  59. 


Volume 

Volume 

Volume 

Volume 

Grams  ol 

of  solu- 

of solu- 

Grams of  of  solu- 

of  solu- 

Solutions of  the 

H2O  per 

tion 

tion 

Solutions  of  the 

H2O  per      tion 

tion 

salts. 

c.c.  of  the 

cont. 

cont. 

salts. 

c.c.  of  the    cont. 

cont. 

solution. 

1  c.c. 

30  c.c. 

solution. 

1  c.c. 

30  c.c. 

H2O. 

H2O. 

H2O. 

H2O. 

N      NaNOi  

0.9696 

1.031 

30.93 

N/2  Bad.  

0.9842 

1.016 

30.48 

N      NaBr  . 

0  9728 

1  028 

30  84 

N/2  Mg(NO»)» 

0  9789 

1  022 

30  66 

N      KBr  

0  9642 

1  037 

31   11 

N/4  NaNO, 

0  9881 

1  012 

30  36 

N      BaCl2  

0  9700 

1  031 

30  93 

N  /4  NaBr 

0  9922 

1  008 

30.24 

N      Mg(NO3)2.... 

0.9527 

1  049 

31  48 

N  /4  KBr  

0  9868 

1.013 

30.39 

N/2  NaNO,  

0.9803 

1.020 

30.60 

N/4  BaCU  

0.9917 

1.008 

30.24 

N/2  NaBr 

0  9852 

1  015 

30  45 

N/4  Me(NOa)i 

0  9877 

1  012 

30  36 

N/2  KBr.. 

0.9804 

1.020 

30.60 

Action  of  Hydrated  and  Nonhydrated  Salts  on  Saponification. 


97 


TABLE  60. 


No. 

Solutions 

N. 

15° 

24  hr*. 

25° 
24  hrs. 

35° 
24  hrs. 

25° 

48  hrs. 

Solutions 

N/2. 

25° 
24  hrs. 

35° 
24  hrs. 

Solutions 

N/4. 

25° 
24  hrs. 

1 

2 

f  30  c.c.      1 
I     H20.      / 

0.086 
0.086 

0.155 
0.154 

0.394 
0.395 

0.296 
0.297 

f   30  c.c.     \ 
I    H80.      I 

0.153 
0.155 

0.394 
0.393 

(  30c.c.     \ 
1    H,0.      ] 

0.153 
0.154 

3 

4 

/30.93c.c.  1 
\NaNOj.    } 

0.105 
0.105 

0.225 
0.224 

0.692 
0.691 

0.490 
0.488 

/30.60c.c.  \ 
[NaNOi.    / 

0.223 
0.222 

0.663 
0.661 

/30.36C.C.  1 
\NaNO,.   / 

0.210 
0.211 

5 
6 

f30.84c.c.  1 
\   NaBr.     / 

0.104 
0.105 

0.228 
0.227 

0.703 
0.704 

0.517 
0.517 

f30.45c.c.  \ 
1   NaBr.     / 

0.222 
0.223 

0.660 
0.661 

/30.24  c.c.  \ 
1   NaBr.     j 

0.210 
0.209 

7 
8 

f31.ll  c.c.  1 
\     KBr.       / 

0.107 
0.106 

0.230 
0.231 

0.690 
0.689 

0.507 
0.505 

f30.60c.c-  1 
[    KBr.      } 

0.224 
0.224 

0.667 
0.668 

f30.39c.c.  1 
1    KBr.       ] 

0.216 
0.217 

9 
10 

/30.93c.c.  1 
\   BaCl2.     / 

0.142 
0.144 

0.377 
0.377 

1.38 
1.37 

0.967 
0.965 

f30.48c.c.  1 
1  Bad,.     / 

0.315 
0.316 

1.09 
1.09 

/30.24  c.c.  1 
1  BaCl2.     / 

0.267 
0.268 

11 
12 

/31.48c.c.  "1 
\Mg(NO,),/ 

0.125 
0.126 

0.326 
0.325 

1.35 
1.36 

0.883 
0.885 

/30.66c.c.  1;  0.285 
lMg(NO,),/   0.286 

1.13 
1.14 

f30.36c.c-  1 
\Mg(NO,)J 

0.261 
0.263 

TABLE  61. 


|  Volume 

Volume 

, 

apSi-d 

Volume  Volume 

Grams  of 

of  solu- 

of solu- 

• 

Grams  of 

of  solu-  of  solu- 

Solutions of  the 

HjO  per 

tion 

tion 

Solutions  of  the 

H2O  per 

tion 

tion 

salts.               [c.c.  of  the 

cont. 

cont. 

salts. 

c.c.  of  the 

cont. 

cont. 

solution. 

1  c.c. 

30  c.c. 

solution. 

1  c.c. 

30  c.c. 

H2O. 

H2O. 

H2O. 

H2O. 

N      KI  

0.9513 

1.051 

31.53 

N/2  MgSO«  

0.9950 

.005 

30.15 

N      Nal 

0  9611 

1  040 

31   20 

N  /2  Sr(NOj)»  .    . 

0  9758 

.025 

30.75 

N      CaBr2 

0  9597 

1  042     31   2fi 

N/4  KI 

0  9864 

014 

30.42 

N      MgSO4 

0  9938 

1  006 

30  18 

N  /4  Nal  

0  9904 

.010 

30.30 

N      Sr(NOj)2             0  9509 

1  052 

31  56 

N  /4  CaBrj                    0  9891 

Oil 

30.33 

N/2  KI                     '0  074fi 

1  026 

30  78 

N/4  MgSO<  ''  0.9958 

.004 

30.12 

N/2  Nal  

0.9809 

1.019 

30.57 

N/4  Sr(NOj)2  0.9881 

.012 

30.36 

N  /2  CaBrs  .  .  . 

0.9825 

1.018 

30.54 

TABLE  62. 


No. 

Solutions 

N. 

15° 
24  hrs. 

25° 
24  hrs. 

35° 
24  hrs. 

25° 
48  hrs. 

Solutions 

N/2. 

25° 
24  hrs. 

35° 
24  hra. 

Solutions 

N/4. 

25° 
24  hrs. 

1 

f  30  c.c.    1 

0.086 

0.155 

0.395 

0.296 

<  30  c.c.    \ 

0.155 

0.395 

f  30  c.c.    \ 

0.155 

2 

I   H20.      / 

0.086 

0.155 

0.394 

0.296 

I   HsO.      / 

0.155 

0.394 

1   H20.     / 

0.154 

3 

/31.53  c.c.l 

0.030 

0.038 

0.057 

0.048 

/30.78  c.c.\ 

0.073 

0.172  i 

/30.42  c.c.l 

0.136 

4 

1     KI.      ) 

0.032 

0.036 

0.056 

0.049 

I     KI.      } 

0.073 

0.172  ' 

I     KI.      ) 

0.135 

5 

/31.20  c.c.l 

0.075 

0.119 

0.341 

0.239 

/30.57  c.c.\ 

0.171 

0.498 

/30.30  c.c.1 

0.176 

6 

1    Nal.      / 

0.077 

0.118 

0.342 

0.241 

1    Nal.      / 

0.172 

0.498 

1    Nal.      / 

0.177 

7 

/31.26  c.c.l 

0.075 

0.134 

0.653 

0.492 

/30.54  c.c.1 

0.153 

0.689 

/30.33  c.c.1 

0.191 

8 

\CaBr2.    / 

0.073 

0.132 

0.651 

0.494 

ICaBrj.    / 

0.151 

0.671 

\CaBrs.    / 

0.192 

9 

f30.  18  c.c.1 

0.095 

0.166 

0.406 

0.346 

/30.15  c.c.\ 

0.150 

0.372 

f  30.12  c.c.\ 

0.138 

10 

lMgSO«.   Jj  0.097 

0.168 

0.408 

0.344 

\MgSO,.   / 

0.148 

0.374 

iMgSO,.   / 

0.140 

11 

f3156c.c.\ 

0.139 

0.297 

1.09 

0.711 

/30.75  c.c-1 

0.287 

0.968 

/30.36  c.c.1 

0.260 

12 

lSr(N03)2/ 

0.139 

0.298 

1.09 

0.713 

lSr(NOs)2/ 

0.285 

0.966 

\Sr(N08)J 

0.261 

Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


TABLE  63. 


Volume 

Volume 

Volume 

Volume 

Grams  of 

of  solu- 

of solu- 

Grams of  of  solu- 

of  solu- 

Solutions of  the     j  H2O  per 

tion 

tion 

Solutions  of  the 

H2O  per 

tion 

tion 

salts.                e.c.ofthe 

cont. 

cont. 

salts. 

e.c.ofthe 

cont. 

cont. 

solution. 

1  c.c. 

30  c.c. 

solution. 

1  c.c. 

30  c.c. 

H2O. 

H2O. 

H20. 

HjO. 

N      LiBr 

0  9739 

1  027 

30  81 

N  /2  SrBr2 

0  9765 

1   024 

30  72 

N       ].i,S(l,  

0  .  9757 

1.025 

30  75 

N  /2  LiNOs   .    . 

0  9825 

1  018 

30  54 

N      BaBr,  

0  .  9529 

1.049 

31.47 

N/4LiBr  

0  9928 

1  007 

30  21 

N      SrBr, 

0  9536 

1  049 

31  47 

N  /4  Li,SO4 

0  9940 

1  006 

30  18 

N      LiNOj  

0  9673 

1  034 

31  02 

N  /4  BaBr2 

0  9871 

1  013 

30  39 

N/2  LiBr  

0.9849 

1.015 

30  45 

N  /4  SrBr,  . 

0  9870 

1   013 

30  39 

N/2  Li,SO4  

0.9882 

1.012 

30.36 

N/4  LiNOj  

0  9891 

1   Oil 

30  33 

N  /2  BaBr,    .  . 

0.9761 

1.025 

30.75 

TABLE  64. 


No. 

Solutions 

N. 

15° 
24  hrs. 

25° 
24  hrs. 

35° 

24  hrs. 

25° 
48  hrs. 

Solutions 

N/2. 

25° 
24  hrs. 

35° 
24  hrs. 

Solutions 
N/4. 

25° 
24  hrs. 

1 
2 

f  30  c.c.    \ 
I   H,0.     ] 

0.087 
0.086 

0.153 
0.153 

0.392 
0.392 

0.297 
0.298 

1  30  c.c.    1 
I   H20.     ] 

0.153 
0.154 

0.392 
0.392 

/  30  c.c.    \ 
1   H20.     / 

0.153 
0.153 

3 

4 

/30.81  0.0.1 
I  LiBr.     / 

0.033 
0.033 

0.053 
0.053 

0.128 
0.129 

0.100 
0.099 

/30.45  c.c.1 
1  LiBr.     J 

0.105 
0.106 

0.307 
0.308 

/30.21  c.c.l 
{  LiBr.     J 

0.161 
0  .  160 

5 
6 

/30.75  c.c.\ 
1  Li,S04.  / 

0.103 
0.105 

0.153 
0.153 

0.306 
0.305 

0.278 
0.277 

/30.36  c.c.l 
1  Li2S04.  / 

0.149 
0.151 

0.297 
0.299 

/30.18  c.e.l 
I  Li,S04.  / 

0.139 
0.139 

7 
8 

/31.47  c.c.1 
\BaBr,.    / 

0.081 
0.081 

0.179 
0.177 

0.705 
0.707 

0.445 
0.446 

/30.75  c.c.\ 
\BaBrj.    J 

0.223 
0.224 

0.813 
0.815 

/30.39  c.c.1 
\BaBr2.    / 

0.241 
0.239 

9 

10 

/31.47  c.c.\ 
1  SrBr,.    / 

0.075 
0.077 

0.096 
0.098 

0.250 
0.248 

0.190 
0.192 

/30.72  c.c.1 
\  SrBr,.    / 

0.130 
0.128 

0.387 
0.385 

F30.39  c.c.\ 
1  SrBr2.    / 

0.164 
0.162 

11 

12 

/31.  02  c.c.1 
\  LiNO8.  / 

0.086 
0.086 

0.1G7 
0.168 

0.508 
0.510 

0.370 
0.371 

/30.54  c.c.\ 
\LiNOt.  J 

0.182 
0.182 

0.555 
0.557 

/30.33  c.c.1 
\LiNOa.  / 

0.201 
0.200 

TABLE  65. 


No. 

Solutions 

N. 

15° 
4  hrs. 

25° 
4  hrs. 

15° 
8  hrs. 

Solutions 

N/2. 

15° 

4  hrs. 

Solutions 

N/4. 

15° 
4  hrs. 

1 

2 

(  30  c.c.   \ 
1    EM).    / 

0.668 
0.667 

3.89 
3.90 

2.58 
2.60 

f  30  c.c.   1 
1    HS0.     ] 

0.680 
0.686 

f  30  c.c.   1 
I    H,0.     1 

0.667 
0.661 

3 
4 

/30.81  c.c.1 
I    LiBr.    ] 

0.309 
0.296 

5.26 
5.24 

2.56 
2.56 

/30.45  c.c.1 
\    LiBr.    / 

0.650 
0.643 

/30.21  c.c.l 
1    LiBr.    / 

0.927 
0.939 

5 
6 

/30.75  c.c.1 
{  Li2S04.  / 

0.742 
0.742 

3.96 
3.97 

3.80 
3.78 

/30.36  c.c.1 
1  Li,S04.  / 

0.667 
0.661 

/30.18  c.c.1 
I  Li2S04.  / 

0.618 
0.605 

7 
8 

/31.47  c.c.1 
\  BaBrj.  / 

0.754 
0.747 

9.27 
9.27 

4.76 
4.76 

/30.75  c.c.l 
i  BaBr,.  ] 

1.050 
1.070 

/30.39  c.c.l 
\  BaBr,.  / 

1.420 
1.430 

9 
10 

/31.47  c.c.\ 
1   SrBr,.   ] 

0.278 
0.279 

3.47 
3.45 

1.78 
1.80 

/30.72  c.c.l 
I   SrBr,.   / 

0.590 
0.588 

/30.39  c.c.1 
1   SrBr,.   / 

0.990 
0.988 

11 
12 

f  3  1.02  c.c.1 
\  LiNO,.  ] 

0.865 
0.865 

6.74 
6.72 

4.44 
4.45 

[30.54  c.c.1 
1  LiNOi.  j 

1.020 
1.020 

/30.33  c.c.l 
\  LiNOj.  J 

1.110 

1.120 

Action  of  Hydrated  and  Nonhydrated  Salts  on  Saponification. 


99 


TABLE  66. 


I 
Volume  Volume 

Volume 

Volume 

Grams  ol 

of  solu- 

of solu- 

Grams of 

of  solu- 

of solu- 

Solutions of  the 

H2O  per 

tion 

tion 

Solutions  of  the 

H2O  per 

tion 

tion 

sails. 

c.c.  of  the 

cont. 

cont. 

salts. 

c.c.  of  the 

cont. 

cont. 

solution. 

1  c.c. 

30  c.c. 

solution. 

1  c.c. 

30  c.c. 

H2O. 

H20. 

H2O. 

HSO. 

N      NaBr 

0  9722 

029 

30  87 

N/2  Ca(NO»)2 

0  9772 

1  023 

30  69 

N      KC1  

0.9704 

030 

30  90 

N/2  MgSO< 

0  9950 

1  005 

30  15 

N      BaClj  

0.9703 

.031 

30.93 

N/4  NaBr  

0  9928 

1  007 

30  21 

N      Ca(NO»)2  .... 

0.9494 

.053 

31.59 

N  /4  KC1  .  . 

0  9912 

1  009 

30  27 

N      MgSO< 

0  9938 

006 

30  18 

N  /4  BaClj 

0  9922 

1  008 

30  24 

N/2  NaBr  

0.9866 

1.014 

30.42 

N/4  Ca(NO$)».  .  . 

0  9883 

1  012 

30  36 

N  /2  KC1  

0  9850 

1  015 

30  45 

N/4  MgSOi 

0  9958 

1  004 

30  12 

N  /2  BaClj 

0.9842 

1.016 

30.48 

TABLE  67. 


No. 

Solutions 

N. 

15° 
4hrs. 

25° 
4hrs. 

15° 
8hrs. 

Solutions 

N/2. 

15° 
4hrs. 

Solutions 

N/4. 

15° 
4hrs. 

1 
2 

/  30  c.c.     \ 
I    H80.      ] 

0.661 
0.667 

3.87 
3.90 

2.58 
2.60 

f  30c.c.     1 
I     H,0.       ] 

0.680 
0.680 

f  30  c.c.     \ 
I    HjO.      / 

0.680 
0.680 

3 

4 

/30.87  c.c.  1 
\   NaBr.     J 

1.05 
1.06 

7.00 
6.98 

4.25 
4.26 

/30.42c.c.  1 
1   NaBr.     / 

0.988 
0.988 

/30.21  c.c.  \ 
\   NaBr.     J 

0.927 
0.929 

5 
6 

f30.90c.c-  \ 
1     KC1.       j 

1.11 
1.12 

7.05 
7.06 

4.33 
4.32 

f30.45c.c.  \ 
\     KC1.       / 

1.11 
1.10 

f30.27c.c.  \ 
I     KC1.       / 

1.05 
1.04 

7 
8 

/30.93c.c.  1 
1   BaCl2.     J 

1.78 
1.76 

13.71 
13.70 

7.23 
7.22 

f  30.48  c.c.  "1 
1  BaCl2.     J 

1.71 
1.69 

f30.24c.c.  1 
1   Bad,.     J 

1.47 
1.45 

9 
10 

f31.59c.c.  \ 
\Ca(NO,),.j 

1.24 
1.25 

9.86 
9.84 

5.25 
5.24 

f30.69c.c.  "I 
lCa(NO,),J 

1.17 
1.19 

f  30.36  c-c.  \ 

lCa(NO,),./ 

1.11 
1.13 

11 
12 

f30.18c.c.  1 
1  MgS04.    J 

1.36 
1.37 

8.50 
8.49 

5.92 
5.94 

f30.15c.c.  1 
\MgS04.    j 

1.10 
1.12 

/30.12c.c.  1 
I  MgS04.    j 

0.926 
0.928 

TABLE  68. 


Volume 

Volume 

Volume 

Volume 

Grams  of 

of  solu- 

of solu- 

Grams of 

of  solu- 

of solu- 

Solutions of  the 

H2O  per 

tion 

tion 

Solutions  of  the 

H2O  per 

tion 

tion 

salts. 

c.c.  of  the 

cont. 

cont. 

salts. 

c.c.  of  the 

cont. 

cont. 

solution. 

1  c.c. 

30  c.c. 

solution. 

1  c.c. 

30  c.c. 

H20. 

H2O. 

H2O. 

H2O. 

N      KI 

0  9508 

1  052 

31  56 

N/2  CaClj 

0  9870 

1  013 

30  39 

N      CaBr2  

0.9599 

1.042 

31  26 

N/2  Mg(NO»)2  

0.9787 

1.022 

30  66 

N      SrCl2   . 

0  9719 

1  029 

30  87 

N/4  KI 

0  9864 

1  014 

30  42 

N      CaCh 

0  9718 

1  029 

30  87 

N/4  CaBr2 

0  9894 

1  Oil 

30  33 

N      Mg(NOt)2... 

0  9529 

1  049 

31  48 

N  /4  SrCl2 

0  9914 

1  009 

30  27 

N/2  KI 

0  9741 

1  027 

30  81 

N/4  CaCl2 

0  9912 

1  009 

30  27 

N  /2  CaBr2  

0.9823 

1  018 

30  54 

N/4  Mg(NOt)a  ... 

0.9877 

1  012 

30  36 

N  /2  SrCl2 

0.9852 

1.015 

30.45 

100          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

The  following  curves  show  the  percentages  of  methyl  acetate  sapon- 
ified at  15°,  25°  and  30°. 


Time 24  hours 

Normal  salt  solutions 
Ester-methyl  acetate 


Temperature 

FlOUBE    7. 


Action  of  Hydrated  and  Nonhydrated  Salts  on  Saponification.         101 

The  following  curves  show  the  amounts  of  methyl  acetate  saponi- 
fied in  24  hours  and  in  48  hours. 


Sr(NOj), 


Tern  perature —  25 
Normal   salt  solutions 
Ester-tnethyl  acetate 


NaCl 


NaBr 


CaBr, 
NaNOj 

LiCI 

KNOs 

BaBr, 


Li  NOj 
MgSO» 

H,0 
Li,  SO. 

Nal 
Sr-Br, 


LiBr 


KI 


24 

Time   in  hours 


48 


FlGUKE    8. 


102 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


The  following  curves  show  the  percentages  of  methyl  acetate  sapon- 
ified, using  normal,  half-normal,  and  quarter-normal  solutions  of  salts. 


0.3 


Time 24  hours 

Temperature,. -2  5° 
Ester-  methyl  acetate 


O.I 


Concentration  of  salt  solutions 
FIGURE  9. 


Action  of  Hydraled  and  Nonhydrated  Salts  on  Saponification.          103 


The  following  curves  show  the  amounts  of  methyl  formate  saponi- 
fied at  the  two  temperatures,  15°  and  25°. 


Time 4  hours 


Normal  salt  solutions 
Ester-methyl  formate 


Tern  perature 
FIOUEB  10. 


CaCIt 


Mg(NOi). 
SrCI, 


BaCU 


BaBr, 

CaBr, 
MgSO, 


KCI 
NaBr 
Li  NO, 


Li8r. 


SrBr, 


Kl 


25' 


104          Conductivities  and  Viscosities  in  Piire  and  in  Mixed  Solvents. 

The  following  curves  show  the  percentages  of  methyl  formate  sapon- 
ified in  4  and  in  8  hours. 


Temperature. —  15 
Normal    salt  solution 
Ester-methyl    formate 


Time  in  hours 

FIGURE    11. 


Action  of  Hydrated  and  Nonhydrated  Salts  on  Saponification.          105 


The  following  curves  show  the  percentages  of  methyl  formate  sapon- 
ified with  normal,  half-normal,  and  quarter-normal  solutions  of  salts. 


Time _4hours 

Temperature. .1 5° 
Ester-methyl  formate 


CaCl, 

MgNOj 
SrCI2 

BaCI, 


Ca(NOj) 

LiNOj 

KCI 

CaBr2 
SrBr, 
NaBr 


LUSO. 


Concentration  of  salt  solutions 
FIQUBE  12. 


106 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


TABLE  69. 


No. 

Solution? 

N. 

15° 
4hrs. 

25° 
4  bra. 

15° 

8hrs. 

Solutions 

N/2. 

15° 
4  hrs. 

Solutions 
N/4. 

15° 
4  hrs. 

1 

2 

f     30  c.c.     \ 
I    H,0.      / 

0.670 
0.667 

3.89 
3.91 

2.60 
2.58 

f    30  c.c.     1 
I    H20.      / 

0.661 
0.667 

f    30  c.c.     \ 
I     H20.      | 

0.680 
0.680 

3 
4 

(31.  56  c.c.  \ 
1      KI.       I 

0.192 
0.188 

2.14 
2.14 

1.09 
1.08 

f.30.81  c.c.  1 
1      KI.        ] 

0.494 
0.501 

[30.42  c.c.  1 
I      KI.        / 

0.667 
0.668 

5 
6 

/31.26c.c.  \ 
{  CaBr2.     j 

0.618 
0.618 

8.59 
8.61 

4.73 
4.71 

f  30.54  c.c.  \ 
1  CaBr2.     / 

0.982 
0.984 

f30.33c.c-  \ 
1  CaBr2.     1 

1.03 
1.05 

7 
8 

/30.87c.c.  1 
1    SrCl2.      / 

2.55 
2.56 

17.44 
17.42 

10.69 
10.71 

/30.45c.c.  1 
I    SrCU.      / 

1.98 
1.98 

[30.27  c.c.  1 
1    SrCl2.      / 

1.64 
1.63 

9 
10 

[30.87  c.c.  1 
\  CaCl2.     J 

2.97 
2.97 

19.86 
19.84 

12.74 
12.76 

[30.39  c.c.  1 
1  CaCl2.     / 

2.18 
2.19 

[30.27  c.c.  1 
I   CaCU.     / 

1.85 
1.84 

11 

12 

/31.48c.c.   \ 
\Mg(N08)2/ 

2.52 
2.54 

17.92 
17.94 

11.31 
11.33 

f  30.66  c.c.  1 

lMg(NO,)2/ 

2.02 
2.01 

[30.36  c.c.  1 
lMg(N08)2/ 

1.75 
1.74 

DISCUSSION  OF  RESULTS. 

From  figures  7  and  8  the  salts  which  produce  the  greatest  increase  in 
velocity,  and  which  have  the  largest  temperature  coefficients,  are 
magnesium  chloride,  calcium  chloride,  strontium  chloride,  barium  chlo- 
ride, magnesium  nitrate,  calcium  nitrate,  and  strontium  nitrate.  All 
of  these  salts  crystallize  with  water  of  crystallization,  varying  from  six 
to  two  molecules  each.  Next  come  salts  that  do  not  have  water  of 
crystallization,  as  sodium  chloride,  sodium  nitrate,  potassium  chloride, 
and  potassium  bromide.  Along  with  these  salts  we  have  some  that 
have  water  of  crystallization,  as  sodium  bromide,  calcium  bromide, 
lithium  chloride,  lithium  bromide,  lithium  nitrate,  and  magnesium 
sulphate,  which  seem  to  be  exceptions  in  the  light  of  the  action  of  the 
other  salts  that  crystallize  with  water. 

Still  more  inexplicable  are  the  curves  for  lithium  sulphate,  sodium 
iodide,  strontium  bromide,  lithium  bromide,  and  potassium  iodide, 
since  these  curves  are  below  the  curve  for  pure  water. 

Take  figure  9.  As  the  dilution  is  increased,  salts  such  as  magnesium 
chloride  show  a  marked  decrease  in  velocity ;  the  salts  with  no  water  of 
crystallization  show  a  less  decrease,  while  the  apparent  exceptions,  such 
as  calcium  bromide,  strontium  bromide,  etc.,  show  a  marked  increase 
in  velocity,  most  of  their  curves  ending  above  the  curve  for  water. 

Let  us  now  consider  the  salts  of  any  one  metal  (figures  7  and  8),  e.  g., 
calcium.  The  chloride  increases  the  velocity  more  than  the  nitrate, 
the  nitrate  more  than  the  bromide.  With  potassium  and  sodium  the 
chloride  has  more  effect  than  the  nitrate,  the  nitrate  more  than  the 
bromide,  the  bromide  more  than  the  iodide.  Of  the  salts  studied, 
the  sulphates  have  the  least  effect.  There  also  seems  to  be  a  gen- 
eral relation  between  the  metals  themselves,  as  magnesium,  calcium, 


Action  of  Hydrated  and  Nonhydrated  Salts  on  Saponification.         107 

strontium,  and  barium,  although  in  the  series  lithium,  sodium,  and 
potassium,  lithium  seems  to  be  an  exception. 

The  curves  for  the  methyl  formate  are  of  the  same  general  character, 
the  only  difference  being  that  the  curves  are  more  extended,  more  being 
above  the  curve  for  water,  probably  due  to  the  larger  amount  of  saponi- 
fication.  The  curves  for  calcium  bromide,  barium  bromide,  and  mag- 
nesium sulphate  with  methyl  formate  do  not  appear  to  be  exceptions,  as 
they  fall  above  the  curves  for  salts  with  no  water  of  crystallization. 

Kellogg1  found  in  the  case  of  the  halides  of  potassium  that  the  chlo- 
ride increased  the  reaction  the  most,  the  bromide  less  and  the  iodide 
least,  normal  potassium  iodide  actually  decreasing  the  velocity.  This 
apparently  points  to  the  fact  that  the  anion  plays  an  important  part 
in  the  total  effect  of  the  salt  on  the  reaction  velocity.  With  more 
dilute  solutions  of  potassium  iodide  (0.25  N)  the  velocity  of  the  reaction 
is  increased.  Kellogg's  curves  for  strontium  chloride,  calcium  chloride, 
and  barium  chloride  show  that  the  cation  must  also  be  taken  into 
account.  A  possible  explanation  of  the  behavior  of  barium  bromide, 
calcium  bromide,  etc.,  may  be  that  the  total  effect  due  to  these  salts  is 
due  to  a  combination  between  a  cation  tending  to  increase  the  velocity 
and  an  anion  tending  to  decrease  the  velocity. 

Kellogg1  also  finds  that  there  is  a  concentration  for  each  salt  that 
will  produce  the  greatest  increase  in  the  velocity  of  the  reaction.  Still 
further  increase  or  decrease  in  the  concentration  from  this  point 
diminishes  the  velocity  of  the  reaction. 

The  action  of  the  apparent  exceptions  which  manifested  themselves 
in  the  study  of  this  problem,  i.  e.,  lithium  nitrate,  lithium  bromide, 
calcium  bromide,  etc.,  may  be  due  to  the  concentration  of  maximum 
saponification  not  having  been  reached,  since  these  salts  on  dilution  all 
increase  the  velocity.  It  is  altogether  probable  that  had  the  dilution 
been  still  farther  increased,  the  apparent  exceptions  would  have  proved 
not  to  be  exceptions  at  all.  This  seems  to  be  all  the  more  true,  in  that 
the  curves  of  figure  10,  where  larger  percentages  of  methyl  formate  are 
saponified,  show  that  barium  bromide,  calcium  bromide,  and  magnesium 
sulphate  have  greater  effect  than  potassium  chloride  and  sodium 
bromide.  Therefore,  it  is  more  than  probable  that  with  solutions 
so  dilute  as  to  reach  the  dilution  of  maximum  saponification,  and  with 
a  larger  amount  of  ester  saponified,  most  of  the  apparent  exceptions  in 
figures  7  and  8  would  no  longer  be  exceptions,  but  we  would  have  a 
general  relation  between  salts  with  water  of  crystallization  and  salts 
without,  the  former  increasing  the  velocity  to  a  larger  extent  than  the 
latter,  having  a  larger  temperature  coefficient  and  decreasing  more 
with  dilution. 

We  are  not  satisfied  with  the  results  with  strontium  bromide,  and 
hope  in  the  near  future  to  do  more  work  with  this  salt. 

'Journ.  Amer.  Chem.  Soc.,  31,  886  (1909);  35,  396  (1913). 


108         Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

The  position  of  the  curve  for  a  salt,  then,  seems  to  be  a  function  of  its 
water  of  crystallization  supplemented  by  an  effect  due  to  the  ions  it 
forms,  e.  g.,  water  of  crystallization  would  place  the  curve  above  the 

curves  of  salts  that  do  not  have  water  of  crystallization;  the  anion  Cl 

would  place  it  above  a  salt  with  an  anion  NOs  or  Br;  and  the  cation 
++  ++   ++        ++ 

Mg  would  place  it  above  a  salt  with  a  cation  Ca,  Sr,  or  Ba.  This 
general  relation  appears  to  hold  among  the  salts  that  have  no  water  of 
crystallization.  On  figures  8  and  9  the  lithium  salts  do  not  seem  to  fit 
in  with  the  sodium  and  potassium  salts,  but  what  has  been  said  about 
calcium  bromide  and  barium  bromide  applies  also  to  these  salts,  as  is 
shown  by  figures  9,  10,  and  11. 

How  can  we  explain  this  increase  in  the  velocity  of  the  reaction 
caused  by  salts  that  have  water  of  crystallization,  and  the  decrease  in 
their  effect  on  dilution. 

Getman  and  Bassett1  showed  that  the  salts  having  water  of  crystal- 
lization are  in  solution  the  most  hydrated.  Assuming  that  the  effect 
of  the  salts  with  water  of  crystallization  is  due  to  their  being  hydrated, 
let  us  see  what  we  should  expect. 

(1)  According  to  Getman  and  Bassett,  the  chlorides,  nitrates,  etc., 
of  such  salts  as  magnesium,  calcium,  strontium,  and  barium  would 
increase  the  velocity  of  the  reaction  to  very  nearly  the  same  extent, 
and  this  is  much  greater  than  with  the  non-hydrated  salts. 

(2)  As  the  concentration  decreases  the  effect  would  be  lessened, 
since  the  total  combined  water  would  be  less,  the  decrease  being  far 
more  rapid  than  in  the  case  of  non-hydrated  salts. 

(3)  From  the  work  of  Pearce,2  who  showed  that  the  hydrating  power 
of  a  cation  is  inversely  proportional  to  its  atomic  volume,  we  should 
expect  the  curves  for  the  salts  of  magnesium,  calcium,  barium,  and 
strontium  with  a  common  anion,  to  be  similar  to  those  found  by  him, 
i.  e.,  magnesium  salts  have  the  greatest  effect,  then  calcium,  strontium, 
and  barium. 

Let  us  see  if  the  experimental  data  confirm  these  conclusions,  based 
on  the  assumption  that  between  salts  that  have  water  of  crystallization 
and  salts  that  do  not,  the  difference  in  action  is  due  largely  to  the 
hydrates  formed  by  the  salts  with  water  of  crystallization. 

Taking  into  consideration  the  facts  brought  out  earlier,  that  the 
dilution  of  maximum  saponification  for  the  apparent  exceptions  had 
not  been  reached,  and  that  larger  percentages  of  ester  saponified  would 
also  tend  to  make  the  apparent  exceptions  not  real,  we  can  draw  from 
the  curves  the  following  conclusions: 

(1)  Salts  with  water  of  crystallization  increase  the  velocity  of  the 
reaction  much  more  than  salts  without  water  of  crystallization. 

Carnegie  Inst.  Wash.  Pub.  No.  60,  15  (1907).  'Ibid.,  180, 57  (1913). 


Action  of  Hydrated  and  Nonhydrated  Salts  on  Saponification.          109 

(2)  The  effect  of  salts  with  water  of  crystallization  decreases,  on 
dilution,  much  more  than  the  salts  without  water  of  crystallization. 

(3)  Among  the  metals  with  common  anions,  as  magnesium,  calcium, 
strontium,  and  barium  chlorides  or  nitrates,  the  curves  are  arranged 
in  the  order  of  the  decreasing  atomic  weights  of  the  cations. 

We  therefore  conclude  that  between  salts  with  water  of  crystalliza- 
tion and  salts  without,  on  the  saponification  of  esters,  the  difference 
in  action  is  probably  due  to  the  chemical  difference  between  hydrated 
and  nonhydrated  salts  or  between  free  and  combined  water.  As  an 
explanation  of  this  difference  in  action  we  offer  the  suggestion  that 
the  combined  water  is  more  highly  ionized  than  free  water,  and  with 
hydrated  salts  we  have  this  effect  added  to  the  salt  effect  shown  in 
the  case  of  nonhydrated  salts. 

Pearce1  shows  that  the  cations  are  the  ones  that  are  most  strongly 
hydrated,  the  anions,  if  hydrated  at  all,  being  only  slightly  so. 

The  curves  we  find  for  the  halides  of  potassium  show  that  in  the 
saponification  of  esters  the  anions  play  an  important  part.  This  is  in 
line  with  what  Kellogg2  found.  It  therefore  seems  probable  that  the 
anions  are  also  hydrated  to  a  certain  extent. 

But  how  can  the  larger  temperature  coefficient  of  reaction  velocity  of 
the  hydrated  salts  be  accounted  for,  since  with  rise  in  temperature  the 
hydrates  become  less  complex?  A  study  of  figures  8  and  1 1  shows  that, 
temperature  being  constant,  the  amounts  of  ester  saponified  in  the 
presence  of  the  hydrated  salts,  as  the  time  increases,  is  much  greater 
than  the  amounts  saponified  in  the  presence  of  the  non-hydrated  salts. 
This  is  probably  due  to  the  larger  amount  of  acid  formed  by  the 
saponification  of  the  ester  by  combined  water.  With  rise  in  tempera- 
ture the  larger  amount  of  acid  would  give  us  a  larger  temperature 
coefficient.  Another  factor  to  be  taken  into  consideration  is  the 
hydrolysis  of  the  hydrated  salts,  which,  though  negligible  at  low  tem- 
perature, increases  greatly  as  the  temperature  rises.  The  increased 
fluidity  of  the  solution  must  also  be  taken  into  account,  the  hydrates 
becoming  less  complex  with  rise  in  temperature.  These  three  factors 
would  probably  offset  the  decomposition  of  the  hydrates,  giving  less 
combined  water,  with  the  result  that  we  should  have  a  larger  tempera- 
ture coefficient  of  reaction  velocity  for  the  hydrated  than  for  the  non- 
hydrated salts. 

From  the  standpoint  of  hydrates  breaking  down  with  rise  in  tem- 
perature, let  us  consider  the  effect  of  such  rise  on  chemical  reactions 
in  general.  The  influence  of  temperature  on  the  velocity  of  reactions 
is  usually  very  great. 

Berthelot3  showed  that  the  velocity  with  which  an  ester  is  formed  is 
about  22,000  times  as  great  at  200°  as  at  7°. 

'Carnegie  Inst.  Wash.  Pub.  No.  180,  57  (1913). 

Vourn.  Amer.  Chem.  Soc.,  31,  886  (1909);  35,  396  (1909). 

'Essai  de  Mecanique  Chimique,  2,  93  (1879). 


110          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

Spohr1  finds  that  cane  sugar  is  inverted  5  times  as  rapidly  at  55°  as 
at  25°.  Various  attempts  have  been  made  to  explain  the  effect  of  rise 
in  temperature  on  the  velocity  of  chemical  reactions.  A  decrease  in 
the  viscosity  of  the  medium  with  rise  in  temperature,  allowing  the 
ions  to  move  more  rapidly,  a  decrease  in  the  mass  of  the  ion  with  rise 
in  temperature,  and  the  increased  kinetic  energy  of  the  molecules  and 
ions,  have  been  cited  as  the  causes  of  the  great  increase,  with  rise  in 
temperature,  in  reaction  velocity.  None  of  these  theories  seems  suffi- 
cient to  account  for  such  an  increase  in  the  velocity  of  reactions  as  was 
noticed  by  Berthelot  and  others. 

From  the  solvate  theory  of  solution  we  see  that  ions  and  molecules 
attract  to  them  molecules  of  the  solvent.  We  should  expect  these  ions 
to  react  more  slowly,  especially  if  the  solvate  were  very  complex,  since 
the  solvate  would  act  as  a  protective  covering  around  the  ions  and 
diminish  the  velocity  with  which  they  would  react  with  one  another. 
But  suppose  the  temperature  is  raised;  the  solvate  would  become  less 
and  less  complex,  until  finally  the  ions  would  not  have  any  appreciable  pro- 
tective covering.  In  such  a  case,  where  the  ions  are  in  direct  contact, 
we  should  expect  the  velocity  of  the  reaction  to  be  greatly  increased. 

We  therefore  offer  this  tentative  suggestion  as  one  of  the  most 
important  causes  of  the  increase  in  the  velocity  of  chemical  reactions 
with  rise  in  temperature. 

In  conclusion,  we  can  say  that  the  chemical  differences  between  free 
and  combined  water,  brought  out  in  the  study  of  this  problem,  are 
strictly  analogous  to  the  physical  differences  between  free  and  combined 
water,  as  shown  by  their  power  to  absorb  light,  which  led  us  to  inves- 
tigate this  problem. 

SUMMARY. 

The  following  conclusions  have  either  been  confirmed  or  brought 
out  in  this  investigation : 

(1)  The  reaction  involving  the  decomposition  of  an  ester  proceeds 
slowly  at  ordinary  temperatures,  and  therefore  can  be  readily  and 
accurately  studied. 

(2)  The  hydrolysis  of  the  chlorides  and  nitrates,  etc.,  of  calcium,  mag- 
nesium, strontium,  and  barium  employed  in  the  study  of  this  problem 
is  so  small  that  it  can  not  account  for  the  differences  observed. 

(3)  Salts  with  water  of  crystallization  increase  the  velocity  of  the 
saponification  of  an  ester  to  a  greater  extent  than  salts  with  no  water 
of  crystallization. 

(4)  On  dilution,  the  effect  with  salts  having  water  of  crystallization 
decreases  more  rapidly  than  with  salts  without  crystal  water. 

'Zeit.  phys.  Chem..  2,  195  (1888). 


Action  of  Hydrated  and  Nonhydrated  Salts  on  Saponification.        Ill 

(5)  The  curves  for  the  saponification  of  methyl  formate  are  very 
similar  to  those  for  methyl  acetate. 

(6)  The  large  effect  of  salts  with  water  of  crystallization  is  probably 
due,  in  part,  to  their  being  hydrated,  combined  water  being  more  highly 
ionized  than  free  water. 

(7)  The  amount  of  the  saponification  (and  therefore  the  position  of 
the  curve)  seems  to  be  due  to  the  combined  effect  of  both  cation  and 
anion. 

(8)  It  is  probable  that  anions  as  well  as  cations  are  hydrated. 

(9)  The  hydration  of  cations  is  inversely  proportional  to  their 
atomic  volumes. 

(10)  There  seems  to  be  a  dilution  of  maximum  saponification  for 
each  salt. 

(11)  Hydrated  salts  show  a  large  temperature  coefficient,  notwith- 
standing the  decomposition  of  hydrates  with  rise  in  temperature. 

(12)  Decomposition  of  hydrates  may  play  an  important  r61e  in  the 
increased  velocity  of  chemical  reactions  with  rise  in  temperature. 

(13)  The  chemical  differences  between  free  and  combined  water  are 
analogous  to  the  physical  differences. 


CHAPTER  VI. 

EFFECT  OF  NEUTRAL  SALTS  ON  THE  HYDRATION  OF  ACETIC 

ANHYDRIDE. 


By  GERALD  C.  CONNOLLY. 


HYDRATION  OF  ACETIC  ANHYDRIDE. 

The  hydration  of  acetic  anhydride  has  been  studied  by  several 
investigators  with  varying  degrees  of  success. 

Menschutkin  and  Vasilieff,1  in  studying  the  decomposition  of  acetic 
anhydride  by  water,  attempted  to  find  a  constant  for  the  velocity  of 
hydration.  They  did  not  succeed  because  the  two  substances  did  not 
mix  in  all  proportions,  and  no  solvent  could  be  found  that  did  not  act 
either  on  the  substances  themselves  or  on  the  products  of  the  reaction. 
They  studied  the  change  in  a  homogeneous  mixture  of  acetic  acid, 
acetic  anhydride,  and  water,  and  found  that  the  reaction  was  not  as 
rapid  as  had  generally  been  supposed. 

Hinsberg2  showed  that  the  acetic  anhydride  was  soluble  in  the  water 
in  the  anhydride  state,  and  that  the  presence  of  water  was  not  an  obstacle 
to  the  employment  of  the  anhydride  as  such. 

A.  and  L.  Lumiere  and  H.  Barbier3  showed  that  a  solution  of  acetic 
anhydride  in  water  possesses  practically  all  the  properties  of  acetic 
anhydride  and  was  sufficiently  stable  for  acetylation  purposes. 

In  a  second  paper4  A.  and  L.  Lumiere  and  H.  Barbier  stated  that 
12  per  cent  acetic  anhydride  was  soluble  in  water,  solution  taking  place 
immediately  on  shaking.  They  prepared  5  and  10  per  cent  solutions 
of  acetic  anhydride  in  cold  water  and  set  them  aside.  From  these 
solutions  equal  aliquot  parts  were  withdrawn  every  10  minutes  and 
added  to  a  slight  excess  of  aniline,  the  excess  being  known.  Reaction 
took  place  quantitatively  between  the  aniline  and  the  acetic  anhydride 
not  hydrated  by  the  water,  forming  acetanilide  and  an  equivalent  of 
acetic  acid.  The  total  acid  present  was  then  determined  by  titration 
with  a  normal  solution  of  sodium  hydroxide  in  the  presence  of  phenol- 
phthalein.  They  found  that  the  rate  of  hydrations  was  at  first  rapid 
and  then  decreased,  the  rate  being  more  rapid  the  greater  the  initial 
dilution  of  the  anhydride  and  the  higher  the  temperature.  They 
carried  out  two  experiments  at  0°  and  two  at  15°.  They  also  prepared 
alcoholic  solutions  of  the  anhydride  and  showed  that  when  molecular  pro- 
portions were  used,  esterification  was  incomplete,  even  after  a  month. 

Benrath,8  by  means  of  change  in  density,  attempted  to  measure  the 
rate  at  which  the  anhydride  combined  with  the  water  in  a  solution  of 

'Journ.  Russ.  Phys.  Chem.  Soc.,  21, 192  (1889).  'Ibid.,  3,  35,625(1906). 

JBer.  d.  deutsch.  chem.  Gesell.,23,2962  (1890).  'Zeit.phys.Chem.,67,501  (1909). 

'Bull.  Soc.  Chim.,3,  33,783(1905). 

112 


Effect  of  Salts  mi  Hydration  of  Acetic  Anhydride.  113 

acetic  acid.  He  concluded  that,  with  equivalent  quantities  of  anhy- 
dride and  water,  the  reaction  was  mono-molecular. 

Rivett  and  Sidgwick,1  using  dilute  aqueous  solutions,  followed  the 
hydration  by  measuring  the  electrical  conductivity.  The  measure- 
ments were  made  at  25°.  The  velocity  of  hydration  showed  that  the 
reaction  was  mono-molecular  and  was  not  catalyzed  by  hydrogen  ions; 
and  that  beyond  a  certain  point  the  constant  decreases  steadily  with 
increasing  concentration. 

Orton  and  M.  Jones2  concluded  that  the  hydration  of  acetic  anhy- 
dride in  acetic  acid  as  a  solvent  is  a  slow  reaction  of  the  second  order; 
that  on  dilution  the  increase  in  velocity  was  approximately  proportional 
to  the  amount  of  water  present,  and  that  the  relation  of  the  velocity 
factor  to  the  temperature  was  normal.  The  effect  of  catalysts  was  also 
studied.  It  was  found  that  acids  were  powerful  catalysts  of  the  hydra- 
tion. The  effect  was  most  obvious  in  media  containing  but  little  water, 
diminishing  as  the  proportion  of  the  water  increased,  being  least  obvious 
in  pure  water.  The  value  of  the  velocity  factor  was  a  linear  function 
of  the  concentration  of  the  acid.  Alkalies  and  hydrolyzed  salts  were 
also  found  to  act  as  powerful  catalysts  to  the  hydration  in  aqueous 
solutions. 

Philip3  made  a  study  of  the  reaction  between  acetic  anhydride  and 
water  in  glacial  acetic  acid,  by  determining  the  freezing-points  of  the 
mixture  taken  at  frequent  intervals. 

Wilson  and  Sidgwick4  studied  the  rate  of  hydration  of  a  number  of 
acid  anhydrides.  The  rate  of  formation  of  the  acid  was  determined 
by  measuring  the  increase  in  the  electrical  conductivity  of  a  solution 
of  the  anhydride  in  water. 

NEUTRAL  SALT  ACTION. 

There  has  been  but  little  experimental  work  done  on  the  problem  of 
the  effect  of  the  neutral  salt  on  the  hydrolysis  by  water  alone.  Most 
of  the  work  has  had  to  do  with  the  effect  exerted  by  a  neutral  salt  on  the 
activity  of  an  acid  used  to  catalyze  the  reaction.  This  latter  is  what 
we  generally  understand  by  the  term  "neutral  salt  action."  It  has 
been  studied  mainly  in  two  reactions:  (1)  the  rate  of  inversion  of  cane 
sugar  in  the  presence  of  acids,  and  (2)  the  catalytic  hydrolysis  of  esters. 
In  general  it  has  been  found  that  neutral  salt  action  is  not  greatly 
influenced  by  temperature,  and  that  the  influence  of  neutral  salts  is 
regarded  as  independent  of  the  acid  employed  as  catalyst. 

Practically  the  only  work  done  on  the  effect  of  neutral  salts  on  hydra- 
tion alone  is  that  due  to  Kellogg,  published  in  a  series  of  three  articles.5 
He  showed  that  the  rate  of  hydrolysis  of  ethyl  acetate  by  water  is 

'Journ.  Chem.  Soc.,  97,  733,  1677  (1910).         Mourn.  Chem.  Soc.,  103,  1959  (1913). 

'Ibid.,  101, 1708(1912).  'Ibid.,  31, 403  (1909) ;  886(1909);  35,396(1913). 

'Proc.Chem.  Soc.,28,  259. 


114          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

greatly  accelerated  by  potassium  chloride,  bromide,  and  iodide;  also 
by  the  chlorides  of  sodium,  lithium,  calcium,  strontium,  barium,  and 
cadmium.  The  reaction  was  studied  at  100°.  Results  show  that  the 
accelerating  effect  of  lithium  chloride  is  greater  than  that  of  sodium 
chloride,  although  the  degree  of  ionization  is  less,  and  that  the  chlorides 
of  calcium,  barium,  and  strontium  have  a  greater  effect  than  either 
sodium  or  potassium  chloride,  although  they  are  less  ionized.  Cadmium 
chloride,  the  least  ionized  of  all  the  chlorides  studied,  produced  the 
greatest  effect.  Kellogg  concluded  that  the  effect  produced  by  a 
neutral  salt  on  the  hydrolysis  of  ethyl  acetate,  is  due  to  a  specific  influ- 
ence on  the  non-ionized  portion  of  the  salt  rather  than  to  any  function 
of  the  ions. 

STATEMENT  OF  THE  PROBLEM. 

Jones  and  Anderson1  found  that  the  absorption  spectra  of  salts  like 
neodymium  chloride  and  nitrate,  when  dissolved  in  non-absorbing 
solvents  such  as  water  and  alcohol,  depended  largely  on  the  nature  of 
the  solvent  in  which  the  salt  was  dissolved;  e.  g.,  neodymium  chloride 
dissolved  in  water  had  a  different  absorption  spectra  from  that  of 
neodymium  chloride  dissolved  in  methyl  or  ethyl  alcohol.  They  found 
for  the  first  time  what  they  called  "solvent"  bands,  showing  that  the 
dissolved  substance  was  combined  with  more  or  less  of  the  solvent, 
forming  in  the  one  case  "hydrates"  and  in  the  other  "alcoholates." 
Alcoholates  had,  as  would  be  expected,  very  different  resonance  from 
hydrates. 

Jones  and  Strong2  extended  the  work  of  Jones  and  Anderson  to  a 
large  number  of  solvents  and  to  a  fairly  large  number  of  non-absorbing 
salts,  and  showed  that  these  solvents  had  a  marked  influence  on  the 
absorption  spectra  shown  by  salts  dissolved  in  them.  They  were  able 
to  distinguish  between  the  spectra  of  salts  dissolved  in  normal  alcohol 
and  those  in  the  isomeric  alcohol.  It  would  lead  us  too  far  here  to 
discuss  this  work  in  detail;  therefore,  reference  only  can  be  made  to  the 
original  papers. 

Jones  and  Guy3  built  the  most  sensitive  radiomicrometer  constructed 
up  to  that  time,  and  by  means  of  it  they  studied  quantitatively  the 
intensities  of  absorption  lines  and  bands.  They  found  that,  while 
solutions  of  slightly  hydrated  salts  were  about  equally  transparent  with 
pure  water,  solutions  of  strongly  hydrated  salts  were  very  much  more 
transparent  than  pure  water. 

The  work  of  Jones  and  Guy  was  repeated  by  Jones,  Shaeffer,  and 
Paulus,4  using  an  even  more  sensitive  radiomicrometer  constructed 
by  Shaeffer,  and  confirmed  conclusions  reached  by  Jones  and  Guy. 
They  found  solutions  of  strongly  hydrated  salts  which  were  as  much  as 
40  per  cent  more  transparent  than  a  depth  of  pure  water  equal  to  the  water 

'Carnegie  Inst.  Wash.  Pub.  No.  110  (1909).  'Ibid.,  190  (1913). 

'Ibid.,  130  (1910)  and  160  (1911).  'Ibid.,  210  (1915). 


Effect  of  Salts  on  Hydration  of  Acetic  Anhydride.  115 

in  the  solution  in  question.  This  showed  that  water  of  hydration  has  a 
very  different  resonance  from  that  of  pure  water. 

Having  found  this  physical  difference  between  combined  and  free 
water,  the  problem  was  now  to  see  if  there  was  any  chemical  difference. 

Jones  and  Holmes1  studied  the  action  of  strongly  hydrated  salts 
and  slightly  hydrated  salts  on  the  saponification  of  methyl  acetate  and 
of  methyl  formate  in  the  following  manner :  He  measured  the  velocity 
of  saponification  of  the  ester  by  pure  water,  by  a  solution  of  slightly 
hydrated  salts  containing  the  same  amount  of  water  as  the  pure  water 
used,  and  by  solutions  of  strongly  hydrated  salts  containing  the  same 
amount  of  water  as  the  free  water  employed,  and  as  the  water  in  the 
slightly  hydrated  salts.  Taking  into  account  the  hydrolysis  of  the 
strongly  hydrated  salts  he  found  that  these  salts  saponified  much  more 
rapidly  than  pure  water  itself. 

The  reaction  studied  by  Jones  and  Holmes  was  a  very  slow  one  and 
indicated  that  combined  water  has  greater  activity  than  free  water. 
We  wished  to  investigate  the  same  problem,  using  a  reaction  that  pro- 
ceeded much  more  rapidly;  therefore  we  chose  the  reaction  involving 
the  conversion  of  acetic  anhydride  into  acetic  acid. 

EXPERIMENTAL. 
PURIFICATION  OF  MATERIAL. 

Very  pure  acetic  anhydride  was  necessary  for  our  work.  The  phys- 
ical properties  as  described  varied  greatly.  The  boiling-points  given 
ranged  anywhere  from  135°  to  140°,  at  760  mm.  pressure.  The  densi- 
ties given  varied  between  1.07  and  1.09.  From  this  it  can  be  seen  that 
it  was  impossible  to  test  its  purity  by  the  ordinary  simple  means. 
Acetic  acid  is  the  impurity  most  likely  to  be  present  in  the  anhydride, 
and  is  very  difficult  to  detect  if  only  small  amounts  are  present;  0.51 
gram  of  pure  acetic  anhydride,  when  completely  hydrated,  is  equiva- 
lent to  100  c.c.  N/10  solution  of  sodium  hydroxide;  while  the  same 
weight  of  a  mixture  containing  1  per  cent  of  acetic  acid  is  equivalent  to 
99.85  c.c.  This  is  within  the  experimental  error.  The  method  finally 
used  to  purify  the  anhydride  was  to  distill  repeatedly,  using  a  five-bulb 
distilling  head,  discarding  the  first  and  last  fractions.  This  gave  a 
very  pure  anyhdride  with  a  constant  boiling-point. 

The  salts  used  in  this  work  were  the  purest  obtainable.  They  were 
repeatedly  recrystallized. 

APPARATUS  AND  SOLUTIONS. 

All  of  the  glassware  used  was  of  Jena  make.  The  constant  tempera- 
ture baths  were  those  commonly  used  in  this  laboratory.  All  solutions 
were  made  up  gravimetrically,  except  those  of  the  non-hydrated  salts 

'See  Chapter  V. 


116          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

which  were  weighed  directly.  The  solution  of  sodium  hydroxide  used 
in  titration  was  made  up  approximately  half-normal,  from  sodium 
hydroxide  from  alcohol.  It  was  preserved  in  an  apparatus  protected 
from  impurities  in  the  air.  It  was  standardized  by  titrating  against  a 
standard  solution  of  sulphuric  acid  of  about  the  same  strength.  The 
purest  water  obtainable  was  always  used. 

MANIPULATION. 

The  method  in  principle  is  a  modification  of  that  of  Menschutkin  and 
Vasilieff,  and  later  employed  by  A.  and  L.  Lumiere  and  Barbier.  In 
order  that  the  results  should  be  comparable,  the  amount  of  water  pres- 
ent must  be  kept  constant  and  the  specific  gravity  of  the  salt  solution  was 
therefore  first  taken;  this  gave  the  weight  of  1  c.c.  From  analysis, 
that  part  of  the  weight  due  to  the  anhydrous  salt  alone  was  known  for 
each  cubic  centimeter.  This  known  weight  of  salt,  subtracted  from 
the  weight  of  1  c.c.  of  solution,  gave  the  weight  due  to  the  pure  water 
alone.  This,  divided  into  the  weight  of  1  c.c.  of  pure  water  at  that 
temperature,  gave  the  amount  of  solution  in  cubic  centimeters  equiv- 
alent to  1  c.c.  of  pure  water.  The  amount  of  solution  thus  calculated 
was  pipetted  into  a  250  c.c.  Jena  bottle.  An  equivalent  of  100  c.c.  of 
pure  water  was  taken  in  all  determinations.  The  bottle  was  suspended 
in  the  constant-temperature  bath.  There  was  also  placed  in  the  bath 
a  bottle  containing  the  anhydride  and  a  number  of  small  empty  bottles 
of  50  c.c.  capacity. 

When  all  had  come  to  the  temperature  of  the  bath,  the  bottle  was 
removed  and  5  c.c.  of  the  anhydride  introduced.  Time  was  reckoned 
from  when  the  anyhydride  was  first  added.  Solution  took  place  imme- 
diately on  shaking,  except  in  the  case  of  the  very  concentrated  solutions. 
Equal  aliquot  portions  were  removed  and  placed  in  the  small  50  c.c. 
bottles,  the  whole  being  kept  in  the  bath.  These  small  bottles  were 
removed,  first  every  5,  then  every  10  minutes,  and  a  slight  known  excess 
of  aniline  added.  This,  on  shaking,  combines  with  the  residual  acetic 
anhydride,  precipitating  acetanilide  and  liberating  an  equivalent  of 
acetic  acid. 

The  total  amount  of  acetic  acid  was  then  titrated  directly  in  the 
bottle,  using  the  half-normal  solution  of  sodium  hydroxide  in  the  pres- 
ence of  phenolphthalein  as  indicator.  Corallin  had  been  tried,  but 
phenolphthalein  was  more  satisfactory.  The  amount  of  acetic  acid 
due  to  the  water  alone  was  then  calculated,  using  the  simple  formula 
y  =  2z — x,  where  y  is  the  amount  of  acetic  acid  due  to  the  water  alone,  z  is 
the  total  amount  of  acetic  acid  as  measured  on  the  burette,  and  x  is  the 
total  amount  of  acetic  acid  that  can  be  formed  if  all  the  acetic  anhydride 
has  been  completely  hydrated. 

Two  temperatures,  15°  and  25°,  were  employed.  Only  one  concen- 
tration of  acetic  anhydride  was  used  (approximately  5  per  cent), 


Effect  of  Salts  on  Hydration  of  Acetic  Anhydride.  117 

because  if  two  were  employed  the  results  would  not  then  be  comparable 
on  account  of  volume  changes.     For  the  salts  molar  m,  half-molar  •?  > 

and  quarter-molar  -j  solutions  were  taken  in  all  cases,  and  usually 

another  solution  as  concentrated  as  possible.     The  time  is  expressed  in 
minutes  and  all  the  results  in  percentages. 

TABLE  70. — Potassium  chloride. 


Concentration  (temperature,  15°).' 

Concentration  (temperature,  25°). 

HSO. 

3M. 

2M. 

M. 

M/2. 

M/4. 

H,0. 

3M 

2M. 

M. 

M/2. 

M/4. 

min. 

5 

32.86 

15.4 

22.4 

30.00 

32.22 

33.10 

43.77 

25.36 

28.70 

41.43 

42.56 

43.54 

10 

57.69 

30.8 

39.5 

52.15 

54.97 

56.72 

71.93 

46.60 

56.26 

65.57 

69.47 

71.34 

20 

79.95 

50.2 

60.4 

75.22 

77.79 

80.31 

93.06 

72.20 

85.78 

89.07 

91.59 

92.84 

30 

90.15 

64.1 

73.9 

86.52 

88.53 

90.89 

98.65 

86.36 

92.66 

96.15 

97.53 

98.90 

40 

95.17 

74.1 

83.4 

92.09 

94.16 

95.54 

99.91 

93.64 

96.64 

98.10 

99.00 

99.59 

50 

97.44 

81.98 

89.5 

95.16 

96.72 

97.32 

60 

98.34 

86.8 

92.6 

97.58 

98.02 

98.53 

TABLE  71. — Calcium  chloride. 


Time. 

Concentration  (temperature,  15°). 

Concentration  (temperature,  25°). 

H2O. 

4M. 

M. 

M/2. 

M/4. 

H,0. 

4M. 

M. 

M/2. 

M/4. 

mtn. 
5 
10 
20 
30 
40 
60 
60 

32.86 
57.69 
79.95 
90.15 
95.17 
97.44 
98.34 

2.55 
15.60 
41.80 
58.65 
70.55 
78.85 
85.06 

32.54 
56.88 
80.70 
91.87 
96.65 
98.72 
99.84 

32.92 
57.71 
80.97 
91.86 
96.39 
98.35 
99.07 

32.88 
57.24 
81.06 
91.96 
96.82 
98.34 
98.59 

43.77 
71.93 
93.06 
98.65 
99.91 

20.53 
50.25 
78.65 
92.74 
96.70 

50.18 
69.46 
97.32 
100.00 

50.49 
72.81 
95.56 
97.66 
100.00 

49.99 
73.45 
95.92 
99.36 
100.00 

TABLE  72. — Magnesium  Chloride. 


Concentration  (temperature,  15°). 

Concentration  (temperature,  25°). 

Time. 

H,0. 

4M. 

2M. 

M. 

M/2. 

M/4. 

H8O. 

4M. 

2M. 

M. 

M/2. 

M/4. 

min. 

5 

32.86 

0.97 

20.30 

30.80 

31.30 

32.82 

43.77 

1.86 

29.70 

40.00 

41.60 

43.09 

10 

57.69 

13.50 

37.45 

48.30 

51.05 

55.58 

71.93 

25.05 

56.51 

69.41 

70.59 

71.98 

20 

79.95 

24.12 

59.11 

74.82 

77.96 

80.60 

93.06 

72.25 

80.26 

89.79 

92.39 

93.31 

30 

90.15 

50.50 

74.02 

87.39 

89.88 

90.76 

98.65 

93.21 

92.37 

96.06 

97.00 

97.89 

40 

95.17 

67.45 

82.68 

92.25 

93.74 

95.89 

99.91 

96.53 

97.46 

98.45 

98.84 

98.88 

60 

97.44 

84.00 

88.87 

96.42 

97.20 

98.13 

60 

98.34 

88.26 

92.64 

97.26 

98.14 

98.60 

118          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

DISCUSSION. 

There  is  one  difficulty  in  the  study  of  this  problem  that  must  first 
be  pointed  out,  i.  e.,  the  necessary  use  of  a  strong  alkali  (half-normal 
solution  of  sodium  hydroxide)  with  which  to  titrate  the  acetic  acid 
formed.  This  necessarily  introduces  some  error,  since  a  difference  of 
0.1  c.c.  in  reading  the  burette  would  make  a  difference  of  over  1  per 
cent.  A  more  dilute  solution  of  an  alkali  could  not  be  used,  since  too 
large  a  quantity  of  such  a  solution  would  be  required. 

The  rate  of  the  decomposition  of  the  anhydride  is  at  first  very  rapid, 
then  gradually  decreases  as  the  hydration  approaches  completion. 
In  this  respect  the  reaction  differs  from  similar  ones  studied,  such  as 
the  hydrolysis  of  esters,  since  in  these  cases  the  reactions  are  reversible. 
There  is  also  a  marked  increase  in  the  rate  of  decomposition  of  the 
anhydride  with  rise  in  temperature. 

The  values  obtained  for  the  salt  solutions,  having  a  fixed  quantity  of 
water,  are  compared  with  those  obtained  with  pure  water.  All  of  the 
very  concentrated  salt  solutions  show  a  marked  decrease  in  the  rate  of 
decomposition  of  the  anhydride.  This  retardation  is  greater  with  the 
two  hydrated  salts  used  than  with  the  non-hydrated  salt.  However, 
with  the  more  dilute  solutions,  molar,  etc.,  this  is  not  the  case.  Mag- 
nesium chloride  and  potassium  chloride  hinder  the  rate  of  decomposi- 
tion to  about  the  same  extent,  this  being  slightly  greater  with  the 
solutions  of  magnesium  chloride.  Calcium  chloride,  on  the  other  hand, 
shows  a  slightly  accelerating  action. 

CONCLUSION. 

These  relations  are  interesting,  but  as  yet  no  general  or  final  con- 
clusions can  be  drawn.  The  problem  is  to  be  investigated  farther, 
using  a  much  larger  number  of  salts,  both  hydrated  and  non-hydrated. 


CHAPTER  VII. 
CONDUCTIVITY  OF  CERTAIN  ORGANIC  ACIDS  IN  ETHYL  ALCOHOL. 


BY  H.  H.  LLOYD  AND  JOHN  B.  WIESEL. 


During  the  past  six  years  a  fairly  thorough  and  systematic  study  of 
the  conductivity  and  dissociation  of  aqueous  solutions  of  organic  acids, 
as  affected  by  temperature  and  by  dilution,  has  been  in  progress  in 
this  laboratory.1  Very  little  work  had  been  done  upon  solutions  of 
organic  acids  in  absolute  ethyl  alcohol  (as  will  become  evident  from  the 
historical  review  that  follows),  and  it  was  therefore  decided  to  extend 
our  investigations  into  this  field.2  Accordingly,  Wightman,  Wiesel, 
and  Jones  undertook  a  preliminary  investigation  of  the  problem,  worked 
out  a  fairly  satisfactory  method  of  procedure,  and  made  conductivity 
measurements  of  nine  organic  acids.3  The  present  investigation  is  a 
continuation  and  extension  of  their  work. 

HISTORICAL  SKETCH. 

The  property  of  water,  which,  in  the  pure  condition,  conducts  elec- 
tricity only  very  slightly,  to  become  conducting  upon  the  addition  of 
many  compounds  which  of  themselves  are  non-conducting,  was  known 
to  De  la  Rive,  and  was  the  subject  of  interesting  experiments  by  Fara- 
day; and  the  fact  that  solvents  other  than  water  exhibit  the  same 
phenomenon  (of  acquiring  conductivity  upon  the  addition  of  non- 
electrolytes)  has  been  studied  by  a  large  number  of  workers;4  but  to 
Kohlrausch  and  Ostwald  belongs  the  credit  of  having  placed  investi- 
gations in  this  field  once  and  for  all  upon  a  firm  scientific  basis. 

The  contributions  of  Kohlrausch  may  be  briefly  described  as  follows: 
(1)  The  development  of  a  convenient  and  precise  method  for  measuring 
conductivity;6  (2)  the  emphasis  on  and  the  determination  of  the 
extent  to  which  conductivity  is  influenced  by  temperature;  (3)  the 
reduction  of  conductivity  measurements  to  definite,  fixed  units;  (4)  the 
reference  of  conductivity  not  to  weight  per  cent,  as  had  been  previously 
done,  but  to  molecular  volumes  of  solutions;  that  is,  volumes  which 
contain  equivalent  amounts  of  the  dissolved  substances;6  (5)  the 
enunciation  of  the  law  of  the  independent  migration  velocities  of  ions. 

Ostwald  made  the  conductivity  method  of  Kohlrausch  more  com- 
plete and  more  practicable  by  reducing  the  apparatus  and  the  manipu- 

lCarnegie  Inst.  Wash.  Pub.  No.  170,  Part  II  (1912);  No.  210,  Chap.  II  (1915). 
'Amer.  Chem.  Journ.,  44,  156  (1910);  46,  56  (1911);  48,  320,  411  (1912);  50,  1  (1913). 
'Journ.  Amer.  Chem.  Soc.,  36,  2243  (1914).     Carnegie  Inst.  Wash.  Pub.  No.  210,  Chap.  Ill 
(1915). 

4Walden:  Zeit.  phys.  Chem.,  8,  433  ff.  (1891);  46,  103  ff.  (1903). 

'Fogg.  Ann.,  138,  379  (1869);  159,  233  (1876);  Wied.  Ann.,  11,  653  (1880);  26,  161  (1885). 
'Wied.  Ann.,  6,  145  (1879). 

119 


120          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

lation  to  the  simplest  possible  terms.1  He  also  showed  how  the  law  of 
independent  migration  velocities  of  ions  could  be  used  to  calculate  the 
/!„,  values  of  organic  acids  from  those  of  their  sodium  salts.2  In  addition, 
just  after  Arrhenius  on  the  basis  of  his  dissociation  theory,  had  pointed 
out3  the  necessity  of  the  parallelism  between  the  strengths  of  acids  and 
their  conductivities,  it  was  Ostwald  who,  led  to  similar  ideas  by  other 
considerations,  succeeded  in  demonstrating4  the  existence  of  such  a 
relation  and  thereby  made  the  first  contribution  to  the  application  and 
significance  of  conductivity  in  questions  of  chemical  affinity.  Further, 
from  thermodynamical  considerations,  Ostwald  derived  his  "dilution 
law,"5  which  permits  the  calculation  (from  the  molecular  conductivities 
of  any  weakly  dissociated  substance  at  different  dilutions)  of  its  affinity 
constant,  which  is  independent  of  the  degree  of  dilution  and  is  condi- 
tioned only  by  the  nature  of  the  dissolved  substance.  During  his 
classical  investigations  in  this  field6  Ostwald  determined  the  affinity 
constants  of  over  240  organic  acids,  and  on  the  basis  of  these  values 
drew  many  conclusions  as  to  the  composition  and  constitution  of  acids, 
thereby  furnishing  organic  chemistry  with  a  new  and  useful  means  of 
investigating  questions  of  constitution.7 

In  a  word,  whereas  Kohlrausch  made  the  determination  of  conduc- 
tivity possible,  by  developing  a  convenient  and  accurate  method  of 
measuring  it,  Ostwald  pointed  out  the  practical  importance  of  these 
measurements  by  indicating  their  bearing  upon  the  solution  of  other 
chemical  problems. 

The  first  serious  attempt  to  study  the  conductivity  of  alcoholic 
solutions  of  organic  acids  was  made  in  1888  by  Hartwig.7  This  investi- 
gator employed  the  Kohlrausch  method  of  measuring  conductivity,  and 
worked  with  solutions  of  formic,  acetic,  and  butyric  acids  of  varying 
concentrations,  at  temperatures  ranging  from  0°  to  30°.  He  found 
that  conductivity  increased  with  rise  in  temperature,  and  calculated  the 
temperature  coefficients  of  conductivity  by  means  of  the  equation 


No  other  conclusions,  however,  can  be  drawn  from  Hartwig's  work, 
because  he  expressed  his  results  m  terms  of  grams  of  acid  in  100  grams 
of  solution  instead  of  in  terms  of  molecular  concentrations. 

In  1889  Kablukoff8  determined  the  conductivity  of  hydrochloric-acid 
gas  in  absolute  alcohol  and  in  alcohol-water  mixtures  at  25°.  He  found 
the  molecular  conductivity  in  pure  alcohol  to  be  about  one-sixteenth  of 
that  in  pure  water.  In  the  mixed  solvent  the  conductivity  increases 

'Zeit.  phys.  Chem.,  2,  561  (1888).  6Zeit.  phys.  Chem.,  2,  36,  270  (1888). 

J/Wd.,  2,  270  (1888);  3,  170(1889).     Amer.  'Ibid.,  3,  170,  241,  369  (1889). 

Chem.  Journ.,46,66  (1911).  'Wied.  Ann.,  2,  33,  58  (1888). 

•Bih.  t.  k.  Ak.,  1884.  "Zeit.  phys.  Chem.,  4,  429  (1889). 
Mourn,  prakt.  Chem.,  30,  93  (1884). 


Conductiirity  of  Organic  Adds  in  Ethyl  Alcohol.  121 

with  increasing  percentage  of  water,  very  slowly  up  to  about  12  per 
cent  of  water  (by  weight),  and  then  more  rapidly.  In  every  case  the 
molecular  conductivity  increases  slowly  with  dilution. 

Wakeman1  measured  the  conductivity  of  several  organic  acids  and 
their  sodium  salts  (as  well  as  of  hydrochloric  acid,  potassium  iodide, 
potassium  chloride,  and  sodium  chloride)  in  alcohol-water  mixtures 
ranging  from  pure  water  to  50  per  cent  alcohol.  Using  Lenz's  values2 
for  the  relative  migration  velocities  of  the  potassium  and  iodine  ions 
in  solutions  of  potassium  iodide,  he  was  able,  by  means  of  his  con- 
ductivity measurements,  to  calculate  the  relative  migration  velocities 
of  the  chlorine,  hydrogen,  and  sodium  ions  in  the  various  mixed  solvents 
which  he  employed.  He  then  had  all  the  data  necessary  to  apply 
Ostwald's  method3  for  the  determination  of  the  nx  values  of  the  organic 
acids  from  those  of  their  sodium  salts.  From  these  values  it  was  a 
simple  matter  to  calculate  percentage  dissociation,  and,  thence,  by 
means  of  Ostwald's  dilution  law,  affinity  constants. 

The  results  obtained  show  that,  for  a  given  acid  at  a  given  dilution, 
increase  in  the  percentage  of  alcohol  causes  a  slight  decrease  in  dissoci- 
ation and  a  more  rapid  decrease  in  the  affinity  constant  (k).  For  a 
definite  alcohol-water  mixture,  the  value  of  k  decreases  regularly  with 
dilution,  which  seems  to  point  to  the  action  of  an  unknown  influence 
which  decreases  the  dissociation  to  such  an  extent  that  the  formula 

2 

— r-  no  longer  holds.     A  very  interesting  feature  of  Wakeman's 

(l-a)y 

work  is  his  attempt  to  extrapolate  the  value  of  molecular  conductivity 
beyond  50  per  cent  alcohol,  in  the  direction  of  100  per  cent  alcohol;  and 
he  postulates  that  the  conductivity  approaches  zero  as  a  limit. 

Schall4  determined  the  conductivity  of  oxalic,  picric,  and  dichlor- 
acetic  acids  in  alcohol  and  in  alcohol-water  mixtures.  The  molecular 
conductivity  of  these  acids  in  water  is  approximately  the  same,  but 
Schall  found  that  replacement  of  the  water  by  alcohol  decreases 
the  conductivity  to  quite  different  extents.  This  he  attributes 
chiefly  to  change  in  the  degree  of  dissociation.  In  alcohol-water 
mixtures  the  acids  were  found  to  behave  very  differently  from  what 
they  did  in  the  pure  solvents.  Some  appear  to  behave  just  the  opposite 
of  what  might  be  expected;  for  example,  picric  acid  gives  a  much 
higher,  and  each  of  the  other  acids  a  much  lower  conductivity  value  in 
water-alcohol  mixtures  than  in  pure  alcohol. 

In  the  meantime  the  conductivity  method  had  been  applied  to  solu- 
tions of  electrolytes  in  a  variety  of  non-aqueous  solvents,  with  the 
result  that  not  a  single  case  was  on  record  in  which  the  Ostwald  dilution 
law  could  be  said  to  apply,  even  approximately.  In  order  to  test  the 

'Zeit.  phys.  Chem.,  11,  49  (1893). 

'Mem.  de  1'Acad.  de  St.  Petersb.,  VII  series,  30,  9  (1882). 

'Zeit.  phys.  Chem.,  2,  270  (1888);  3,  170  (1889).     See  also  Amer.  Chem.  Journ.,  46,  66  (1911). 

4Zeit.  phys.  Chem.,  14,  701  (1894). 


122          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

validity  of  this  law  for  alcoholic  solutions,  Wildermann,  in  1894,  carried 
out  an  investigation1  on  the  conductivity  of  certain  organic  acids 
(acetic,  monochloracetic,  dichloracetic,  trichloracetic,  and  succinic)  in 
absolute  alcohol  at  18°.  Absolute  alcohol  was  obtained  free  from 
aldehyde  by  treatment  with  silver  nitrate,  and  from  water  by  heating 
with  lime.  Great  care  was  exercised  in  protecting  it  from  the  air,  and 
a  special  apparatus  was  constructed,  with  the  aid  of  which  a  measured 
quantity  could  be  drawn  out  of  the  supply  bottle  directly  into  the 
conductivity  cells.  Because  of  the  small  values  of  the  conductivities 
to  be  measured,  the  cells  were  constructed  so  as  to  have  small  constants; 
but  in  spite  of  this  the  external  resistances  required  were  high,  and 
hence  a  graphite  resistance  rheostat  was  employed  to  make  the  tone 
minima  more  distinct.  Before  using,  the  cells  were  washed  with  running 
water  for  from  8  to  10  days.  They  could  not  be  dried  in  the  ordinary 
way  by  washing  with  alcohol  and  then  evaporating  the  alcohol,  because  it 
was  found  that  the  alcohol  which  was  in  contact  with  the  platinized  elec- 
trodes was  to  some  extent  oxidized  by  the  air  to  acetic  acid.  Therefore, 
after  being  washed  with  water,  the  cells  were  allowed  to  drain  com- 
pletely, after  which  alcohol  was  introduced  so  as  to  wash  the  glass  walls 
without  coming  in  contact  with  the  electrodes,  and  then  the  alcohol 
was  allowed  to  cover  the  electrodes.  Some  of  the  alcohol  was  then 
drawn  off  and  fresh  alcohol  added,  keeping  the  electrodes  continuously 
covered,  and  this  process  repeated  until  the  conductivity  remained 
constant.  This  often  cost  a  half  day  of  work  and  from  300  to  500  c.c. 
of  good  absolute  alcohol.  All  the  solutions  were  made  up  in  the  cell, 
the  most  concentrated  by  introducing  approximately  the  desired  quan- 
tity of  acid,  and  from  this  the  others  by  repeated  replacement  of  a 
portion  of  the  solution  with  fresh  alcohol.  The  strength  of  each  solu- 
tion was  determined  by  titration  of  the  portion  removed. 

Wildermann  found  that  the  conductivity  of  acetic,  monochloracetic, 
and  succinic  acids  was  so  small  that  the  values  were  unreliable,  and  he 
was  content  to  make  the  qualitative  statement  that  between  the  vol- 
umes 10  and  160  these  substances  have  a  molecular  conductivity  which 
increases  almost  directly  proportional  to  the  volume. 

In  order  to  understand  fully  his  conclusions  in  regard  to  the  stronger 
acids,  it  will  be  necessary  to  consider  briefly  a  few  mathematical  rela- 
tions which  Wildermann  deduced  from  the  dilution  law  of  Ostwald. 

As  Ostwald  showed,  the  dilution  law 

p-OyO'.a.  (i) 

for  weakly  dissociated  compounds,  takes  the  form 

/ig-Pfcl (II) 

•Zeit.  phys.  Chem.,  14,  231  (1894). 


Conductivity  of  Organic  Acids  in  Ethyl  Alcohol.  123 

Assuming  that  a  given  acid  in  alcoholic  solution  obeys  the  dilution  law, 
if  we  increase  the  volume  from  v  to  v\,  we  have,  instead  of  (I)  and  (II) 

Mao(Moo-M,l)"l         ,.  (fl\ 

A 

and 


(no 

Combining  (II)  and  (IF)  we  get 


H,      v       p. 
Likewise,  from  (I)  and  (F)  there  results 


2 

Since  i*  >»,  M,I>  *».,  and  therefore  M°°~Mri  >  1,  it  follows  that  ^  is 

M°O~M«  M» 

always  less  than  —  and  approaches  this  value  as  a  limit.     Therefore,  as 
molecular  conductivity  increases  in  concentrated  solutions,  as  long  as 

2 

the  compound  is  slightly  dissociated,  the  values  of  —  £  found  experimen- 

0s 

tally  will  not  differ  greatly  from  —  but,  with  increasing  dilution  and 

greater  degree  of  dissociation,  ^  will  decrease  in  value  and  differ  more 

M» 

and  more  from  J 

t;. 

As  a  result  of  his  conductivity  measurements  on  di-  and  trichloracetic 
and  /3-resorcylic  acids,  Wildermann  drew  the  following  conclusions: 

"  (1)  For  dichloracetic  acid,  when  v<  10  liters  the  values  of  —  are  less 

•  *^ 

than  -*/—  5  above  10  liters  —  is  greater  than  x/—  ;  at  dilutions  from  800 
M  v  Mt  ™  • 

to  2,000  liters  —  becomes  ahnost  equal  to  —  .     The  increase  of  —  is, 

ju>  Mt, 

therefore,  continuous,  not  only  in  concentrated  solutions,  but  also  in 

the  more  dilute,  to  which  the  equation  /i°°  M°°  2M"  —  =  k  should  apply." 

"  (2)  The  same  conclusions  are  even  more  nearly  true  in  the  case  of 

/3-resorcylic  acid,  where  the  increase  of  —  is  almost  proportional  to  the 

/*t 

volume." 

it.  phys.  Chem..  2,  270  (1888). 


124          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

(3)  In  the  case  of  trichloracetic_acid,  with  increase  in  dilution  from 

20  to  300  liters  —  approaches  »|£l;  and  beyond  300  liters  a  relation 
Mf  M  v 

obtains  which  can  be  expressed  approximately  by  the  dilution   law. 

The  fact  that  the  values  of  —  for  solutions  more  concentrated  than 

H» 

300  liters  are  not  in  keeping  with  the  dilution  law,  was  confirmed  in  a 
later  investigation1  by  Wildermann  using  an  independent  method.  An 
explanation  of  this  anomalous  condition  was  offered  in  a  previous 
paper,2  and  will  not  be  discussed  here. 

(4)  It  had  been  pointed  out  by  Ostwald,  that  in  aqueous  solutions 
the  degree  of  dissociation  of  the  same  dilution  of  trichloracetic,  dichlor- 
acetic,  monochloracetic,  and  acetic  acids  showed  a  decrease  in  the 
order  named.     A  like  succession  was  observed  by  Wildermann  for  the 
same  acids  in  alcohol. 

In  summing  up,  Wildermann  says  that  it  is  possible  to  apply  the 
Kohlrausch  method  to  the  determination  of  the  conductivity  of  strong 
organic  and  inorganic  acids  in  alcoholic  solution,  but  that  reliable 
results  could  not  be  obtained  for  such  weak  acids  as  acetic,  monochlor- 
acetic, and  succinic.  However,  even  in  the  cases  where  the  method  is 
best  applicable,  much  time  and  patience  on  the  part  of  the  experimenter 
are  required  to  obtain  results  that  are  at  all  reliable. 

Because  of  the  difficulty  which  he  experienced  in  applying  the  Kohl- 
rausch method  to  the  study  of  the  conductivity  of  weakly  dissociated 
organic  acids  in  absolute  alcohol  as  a  solvent,  Wildermann  employed 
another  method1  for  the  investigation  of  these  acids.  This  consists 
in  the  use  of  a  direct  current  of  high  voltage  (140  to  150  volts).  Polari- 
zation can  be  neglected,  and  since  the  external  resistances  are  large,  the 
measurement  of  conductivity  resolves  itself  into  the  measurement  of  the 
strength  of  the  current.  This  is  done  comparatively  by  means  of  a 
reflecting  galvanometer,  the  solutions  being  contained  in  capillary  tubes 
of  varying  lengths  and  diameters.  The  same  conclusions  were  reached 
as  before,  except  that  the  galvanometer  method  was  found  susceptible 
of  more  general  application  than  that  of  Kohlrausch. 

This  method  has  been  found  open  to  objection  by  Malmstrom,3  who 
prefers  a  method  described  by  Nernst,4  whenever  small  conductivities 
are  to  be  measured. 

In  1902  Hantzsch  and  Voegelen5  found  the  conductivity  method 
capable  of  distinguishing  between  a  true  acid  and  a  so-called  "pseudo 
acid"  which  in  aqueous  solution  partly  breaks  down  into  the  constituent 
ions  of  the  true  acid.  The  application  of  the  method  in  this  instance 
depends  upon  the  different  behavior  of  the  molecular  conductivities  of 

'Zeit.  phys.  Chem.,  14,  247  (1894).  *Ibid.,  14,  622  (1894). 

2Ber.d.deutsch.chem.Gesell.,26, 1782  (1893).     'Ber. d. deutsch.  chem.  Gesell.,  35, 1001  (1902). 
3Zeit.  phys.  Chem.,  22,  331  (1897). 


Conductivity  of  Organic  Acids  in  Ethyl  Alcohol.  125 

true  and  pseudo  acids  in  aqueous  alcoholic  solutions  containing  different 
percentages  of  alcohol.  If  it  is  desired  to  test  whether  a  given  hydrogen 
compound  is  a  pseudo  acid  or  not,  a  true  acid  with  approximately  the 
same  affinity  constant  is  selected,  and  the  molecular  conductivity  of 
both  is  determined  in  aqueous  alcohol  at  a  given  dilution.  Then, 
without  varying  the  dilution,  the  percentage  of  alcohol  is  increased  and 
the  resulting  conductivities  determined.  If  the  molecular  conductivity 
of  the  hydrogen  compound  decreases  much  more  slowly  than  that  of 
the  true  acid,  it  is  a  pseudo  acid.  Hantzsch  and  Voegelen  offered  no 
explanation  for  this  difference  in  behavior,  but  they  supported  their 
discovery  with  widely  different  examples. 

In  the  last  few  years  Heinrich  Goldschmidt1  has  applied  the  con- 
ductivity method  to  alcoholic  solutions  of  organic  acids,  in  order  to 
determine  the  equilibrium  hydrogen  ion-alcohol-water,  and  he  has 
succeeded  in  throwing  new  light  on  the  mechanism  of  esterification  and 
on  the  formation  of  complex  ions. 

In  his  earlier  work  (1895)  on  the  kinetics  of  esterification  in  alcoholic 
solutions  of  organic  acids,  Goldschmidt  had  observed  that  the  catalytic 
action  of  strong  acids  in  accelerating  esterification  was  greatly  dimin- 
ished by  the  addition  of  water.  He  explained  this  on  the  assumption 
that  the  accelerating  action  of  a  strong  acid  was  due  to  the  formation 
of  a  complex  molecule  by  union  of  the  acid  with  alcohol,  and  that  the 
retardation  was  brought  about  by  the  decomposition  of  this  molecule 
by  water.  For  example,  if  hydrochloric  acid  were  acting  as  the  cata- 
lyzer, a  compound  C2H5OH,  HC1  would  be  formed,  and  this  would  be 
acted  upon  by  water  according  to  the  equation 

C2H5OH,  HC1+H20  ^=±  C2H8OH+H2O,  HC1  (I) 

For  a  time2  Goldschmidt  thought  that  only  the  ionized  portion  of  the 
complex  molecule  took  part  in  the  esterification,  and  from  measure- 
ments of  reaction  velocity  he  undertook  to  determine  the  equilibrium 
in  the  reaction 

(C2H5OH,  H+)  +H20  ^=2  C2H5OH+ (H2O,  H+)  (II) 

He  found  that  a  number  of  organic  acids  which  were  esterified  in  the 
presence  of  hydrochloric  acid  showed  the  same  retardation  when  equal 
volumes  of  water  were  added.  From  this  he  concluded  that  his 
assumption  was  correct,  and  for  the  equilibrium 

C(C,H.OH,  H+)  X  CH.O 

C(H,O.H+) 
he  calculated  the  equilibrium  constant  r  to  be  0.15. 

'Zeit.  Elektrochem.,  15,  4  (1909) ;  Zeit.  phys.  Chem.,  70,  627  (1910) ;  Ibid.,  81,  30  (1912) ;  Zeit. 
Elektrochem.,  20,  473  (1914);  Zeit.  phys.  Chem.,  89,  129  (1914). 
2Zeit.  phya.  Chem.,  60,  728  (1907). 


126          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

In  1909  Goldschmidt  noticed  that  when  weaker  acids,  such  as  tri- 
chlorbutyric,  picric,  etc.,  were  used  as  accelerators,  the  effect  of  water 
on  the  reaction  velocity  could  not  be  calculated  by  means  of  the 
constant  r  =  0.15,  but  that  a  much  smaller  value  had  to  be  selected. 
Under  these  conditions  not  only  the  ionized  part  of  the  double  com- 
pound alcohol-acid  was  acting  as  catalyzer  (according  to  equation  II), 
but  also  the  undissociated  portion  (as  represented  in  equation  I). 
Consequently,  the  constant  0.15  was  not  wholly  dependent  upon  the 
relation  between  the  ions,  and  could  not  be  used  to  characterize  the 
equilibrium  hydrogen  ion-alcohol-water.  Goldschmidt,  therefore,  de- 
vised a  new  way  of  obtaining  this  equilibrium  constant,  namely,  the 
determination  of  the  limiting  conductivities  of  acids  in  absolute  and 
aqueous  alcoholic  solutions. 

The  molecular  conductivity  of  an  acid  in  absolute  alcohol  at  infinite 
dilution  is  expressed  by  the  well-known  equation 

l*o=U  +  V  (I) 

where  u  equals  the  migration  velocity  of  the  cation  and  v  that  of  the 
anion.  (It  is  assumed  that  there  are  at  least  two  kinds  of  cations  in  the 
solution — free  hydrogen  ions  and  the  complex  ions  (C2H5OH,  H)+; 
but  since  the  ratio  of  these  two  must  always  be  constant  they  may  be 
considered  as  one.)  As  soon  as  water  is  added,  a  new  kind  of  ion  is 
formed — (H2O,  H)+.  If  the  concentration  of  the  water  added  be 
i]  mols.  per  liter,  and  the  migration  velocity  of  the  ions  (H20,  H)+ 
be  u',  the  conductivity  at  infinite  dilution,  /*„,  will  be  expressed  by  the 
equation 

M»=xM+(l-xX+w  (II) 

where  x  represents  the  fraction  of  a  gram  cation  which  is  still  present  in 
the  original  condition  (as  free  or  alcoholated  hydrogen  ion),  and  1  —a: 
the  fraction  transformed  into  the  hydrated  ion  (H2O,  H.)+  Therefore, 
the  equilibrium  equation  which  expresses  the  distribution  of  hydrogen 
ions  between  alcohol  and  water,  is 


whence 

x  —         RUG.  l  —  x = 

•n+r 

Substituting  these  values  in  (II),  we  have 


and,  since  v  =MO— u,  this  reduces  to 

(u-ti')i?  (Ill) 


Conductivity  of  Organic  Acids  in  Ethyl  Alcohol.  127 

If  we  know  fj.0  and  the  values  of  /*„  corresponding  to  two  or  more 
values  of  TJ,  the  equation  can  be  solved  for  r,  the  equilibrium  constant, 
and  for  u—u',  the  difference  between  the  migration  velocity  of  the 
cations  in  anhydrous  alcohol  and  that  of  the  hydrated  ion — (H2O,H)+ 
in  aqueous  alcohol.  Therefore,  in  order  to  test  the  equation,  it  is  only 
necessary  to  determine  the  conductivity  at  infinite  dilution  of  several 
organic  acids  in  pure  alcohol  and  in  two  or  more  alcohol-water  mixtures. 

The  greatest  experimental  difficulty  in  doing  this  lay  in  the  prepara- 
tion of  pure  absolute  alcohol.  At  first1  metallic  calcium  was  used  for 
this  purpose  (von  Winkler's  method),  but  the  alcohol  prepared  in  this 
way,  although  practically  free  from  water,  was  found2  to  contain  con- 
siderable amounts  of  ammonia  (formed  by  the  action  of  water  upon  the 
calcium  nitride  present  in  the  calcium);  and  hence  this  method  had 
to  be  abandoned.  After  further  experimentation  a  satisfactory  product 
was  finally  obtained  in  the  following  manner : 

Ordinary  95  per  cent  alcohol  was  allowed  to  stand  in  contact  with 
lime  for  some  time  and  then  distilled  (Kailan).  This  treatment 
reduced  the  water  content  to  about  0.006-normal,  or  0.12  gram  per  liter. 
The  same  process  was  repeated,  this  time  using  a  distilling  vessel  and 
condenser  of  copper;  and  the  amount  of  water  present  was  decreased 
to  0.003-normal,  or  0.06  gram  per  liter.  If  a  completely  anhydrous 
alcohol  was  desired,  the  product  of  the  first  distillation  was  treated 
with  calcium.  For  this  purpose  calcium  bars  were  turned  on  a  lathe 
to  remove  the  coating  of  hydroxide,  and  the  bright  metal  was  cut  into 
pieces  the  size  of  a  pea.  An  amount  of  calcium  equal  to  ten  tunes 
the  amount  of  water  present  (about  0.1  gram)  was  introduced  into  the 
alcohol.  The  whole  was  then  heated  for  several  hours  with  a  reflux 
condenser  attached,  and  a  rapid  stream  of  dry  air  was  passed  through 
the  distillation  chamber  to  remove  traces  of  ammonia.  In  this  way 
absolute  alcohol  was  obtained  having  a  specific  conductivity  of  2X  10~7. 

With  a  good  quality  of  alcohol  in  his  possession  Goldschmidt  was 
ready  to  test  his  equation,  and  he  began  his  study  with  hydrochloric 
acid.  First,  the  molecular  conductivity  in  absolute  alcohol  was 
measured  at  dilutions  ranging  from  10  to  5,120  liters;  and  from  the 
values  obtained  for  the  more  dilute  solutions  /u0,  the  conductivity  at 
infinite  dilution  was  calculated  with  the  aid  of  Kohlrausch's  formula 


The  mean  value  of  /zo  proved  to  be  89.  Then  the  conductivity  in 
aqueous  alcohol  containing  different  amounts  of  water  (17)  was  deter- 
mined over  the  same  range  of  dilution,  and  the  limiting  conductivity, 
ju,,  corresponding  to  each  alcohol-water  mixture,  was  estimated  as 
before.  The  results  show  that  the  molecular  conductivity  of  solutions 

'Zeit.  Elektrochem.,  15,  4  (1909).  2Zeit.  phys.  Chem.,  81,  30  (1912). 


128 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


of  hydrochloric  acid  in  alcohol  first  decreases  with  increasing  percentage 
of  water,  then  passes  through  a  minimum  at  rj  =  2,  and  finally  begins  to 
increase  again.  px  behaves  in  a  similar  manner. 

Using  the  value  89  for  no,  and  combining  with  it  the  values  of  ju, 
obtained  for  the  different  water  concentrations  (77),  Goldschmidt  solved 
equation  (III)  for  r  and  u—u',  and  obtained  the  constants  0.0583  and 
42.7  respectively.  The  high  value  of  u—u'  shows  that  the  complex 
ion  (H2O,  H)+  travels  much  more  slowly  than  the  free  or  the  alcoholated 
hydrogen  ion.  The  constancy  of  r  and  u—u'  is  evidence  in  favor  of 
the  correctness  of  the  reasoning  by  which  the  equation  was  developed. 

In  order  to  determine  the  limit  of  applicability  of  the  constants  0.0583 
and  42.7,  Goldschmidt  used  these  constants  to  recalculate  the  values 
of  n,.  He  found  satisfactory  agreement  between  the  calculated  values 
and  those  obtained  experimentally,  so  long  as  the  percentage  of  water 
(77)  in  the  alcohol  did  not  exceed  unity.  (Beyond  this  point  there  was 
a  lack  of  agreement,  caused  no  doubt  by  the  influence  of  water  upon 
the  migration  velocities  of  the  ions.)  Thus,  Goldschmidt  showed  that 
by  means  of  the  constants  0.0583  and  42.7  the  effect  of  water  on  the 
conductivity  of  alcoholic  solutions  of  strong  acids  could  be  calculated. 

TABLE  73. 


Acid. 

*«» 
(Nasalt). 

MM 
(acid). 

47 

89 

Sulphosalicylic  
picric  

41 
51 

83 
93 

Trichloracetic  

46 

88 

Trinitrobenzoic  
Trichlorbutyric  
Dichloracetic  .... 

44 
45 
48 

86 
87 
90 

Salicylic  

44 

86 

Desiring  to  adduce  further  evidence  bearing  on  his  theory,  Gold- 
schmidt took  up  a  study  of  the  conductivity  of  sulphosalicylic  acid. 
He  found  the  behavior  of  this  acid  to  be  entirely  analogous  to  that  of 
hydrochloric.  The  molecular  conductivity  in  alcohol  was  influenced 
to  the  same  extent  by  the  addition  of  water,  and  the  amount  of  this 
influence  could  be  calculated  with  the  aid  of  the  constants  0.0583  and 
42.7.  The  limiting  conductivity  in  absolute  alcohol  (MO)  was  estimated 
both  from  the  nx  value  of  the  sodium  salt  (by  the  method  of  Ostwald), 
and  from  the  nx  values  of  the  free  acid  in  aqueous  alcohol  (by  means  of 
equation  III) ,  and  the  figure  obtained  was  the  same  (83)  in  both  cases. 

Satisfied  with  the  results  of  his  investigations  upon  the  stronger 
acids,  Goldschmidt  next  turned  his  attention  to  the  weaker — picric, 
trichloracetic,  trinitrobenzoic,  trichlorbutyric,  dichloracetic,  and  sali- 
cylic. He  observed  that  in  every  case  the  molecular  conductivity 


Conductivity  of  Organic  Acids  in  Ethyl  Alcohol. 


129 


increased  with  increase  in  the  concentration  of  the  water  present  in  the 
solution.  The  nx  values  for  the  weaker  acids  could  not  be  determined 
directly,  as  had  been  done  previously  for  the  strong  acids,  and  Gold- 
schmidt  employed  the  method  of  Ostwald  for  calculating  the  limiting 
conductivity  of  an  acid  from  that  of  its  sodium  salt.  The  values 
obtained  are  given  in  table  73,  which  shows  that  the  difference  in  the 
migration  velocities  of  the  anions  is  very  slight. 

From  the  MM  values  of  the  acids,  the  degrees  of  dissociation  (a)  and 
the  affinity  constants  (k)  were  calculated.  The  values  of  k  given  in 
table  74  were  obtained  for  the  solutions  in  absolute  alcohol. 

TABLE  74. 


Acid. 

k 

1.8  X10-4 

1.5X10-' 

Trinitrobenzoic  
Trichlorbutyric  

8.3X10-' 

i.oxio-7 

7.2  X10-8 

2.4X10-' 

The  influence  of  water  upon  the  values  of  k  was  also  investigated, 
with  the  result  that  k  was  found  to  increase  regularly  with  increasing 
concentration  of  water.  The  law  of  this  increase  was  expressed  in  the 
form  of  an  equation  which  Goldschmidt  deduced  from  theoretical 
considerations. 

In  conclusion,  Goldschmidt  showed  that  the  influence  of  the  weaker 
acids  upon  the  velocity  of  esterification  could  be  calculated  by  means 
of  the  constant  0.0583.  The  calculated  values  agreed  very  satisfac- 
torily with  those  obtained  experimentally. 

This  work  of  Goldschmidt  has  been  discussed  at  some  length,  not 
because  it  is  a  final  word  on  the  subject  with  which  it  deals,  but  for  the 
reason  that  it  presents  an  excellent  illustration  of  the  ever  increasing 
applicability  of  the  conductivity  method  to  the  solution  of  chemical 
problems. 

EXPERIMENTAL  WORK. 
REAGENTS. 

The  alcohol  used  in  this  investigation  was  prepared  in  the  following 
way:  Ordinary  95  per  cent  alcohol  was  heated  for  several  days  with 
lime  in  a  copper  tank  provided  with  a  ground-brass  stopper  and  reflux 
condenser,  and  then  distilled  through  a  block-tin  condenser.  The 
product  of  this  distillation  was  reheated  with  fresh  lime  and  again 
distilled,  the  first  and  last  portions  of  the  distillate  being  discarded. 


130          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

The  receiver  for  the  distillate  was  a  6-liter  Jena  glass  bottle,  having  a 
three-holed  paraffined  cork  as  stopper.  Through  one  hole  passed  a 
siphon,  through  another  an  adapter  with  a  glass  stopcock,  and  through 
the  third  a  calcium  chloride-soda  lime  tube.  In  this  way  the  alcohol 
was  well  protected,  during  distillation,  from  impurities  in  the  air,  and 
small  quantities  sufficient  for  making  up  the  solutions  could  be  drawn 
off  without  exposing  the  main  supply.  By  exercising  extreme  care 
it  is  possible  by  this  method  to  obtain  alcohol  having  a  specific  gravity 
of  0.78506  to  within  the  limit  of  error  ±  0.00002.  According  to  Circular 
19  of  the  Bureau  of  Standards  such  alcohol  is  pure.  The  alcohol 
employed  in  the  conductivity  measurements  varied  in  specific  gravity 
from  0.78506  to  0.78517,  the  latter  containing  99.964  per  cent  alcohol. 

The  organic  acids  were  obtained  from  Kahlbaum  and  Schuchardt. 
The  same  methods  of  purifying  them  were  employed  as  when  the  con- 
ductivities of  these  acids  were  determined  in  aqueous  solution.1  After 
purification  they  were  carefully  dried  in  a  vacuum  desiccator  containing 
sulphuric  acid.  Whenever  practicable  the  melting-points  of  the  acids 
were  taken  as  one  criterion  of  purity. 

Aqueous  solutions  of  ammonia  were  used  for  titration  purposes,  since 
ammonia  had  been  found  by  Wightman,  Wiesel,  and  Jones2  to  give  the 
most  satisfactory  results.  These  solutions  were  prepared  as  follows: 
Concentrated  ammonia  was  heated  and  the  gas  which  was  given  off 
passed  first  over  sodium  hydroxide,  which  collected  a  large  part  of 
the  water-vapor  and  any  carbon  dioxide;  then  over  sodium,  which 
absorbed  the  remainder  of  the  water-vapor;  and  finally  into  a  weighed 
quantity  of  conductivity  water  in  a  measuring  flask,  until  the  amount 
necessary  to  make  a  0.1  N  solution  was  dissolved.  By  diluting  this 
solution  with  conductivity  water  any  desired  strength  of  ammonia 
could  be  obtained.  The  normality  of  each  solution  employed  was 
determined  by  titration  against  standard  sulphuric  acid. 

Corallin  (rosolic  acid)  was  used  as  the  indicator,  because  it  is  sensi- 
tive to  ammonia  and  to  organic  acids,  and  is  not  sensitive  to  carbon 
dioxide  except  when  the  latter  is  present  in  fairly  large  quantity.  The 
end-point  with  corallin  is  not  quite  as  sharp  and  distinct  as  with 
phenolphthalein,  and  considerable  practice  is  necessary  before  reliable 
results  can  be  obtained. 

APPARATUS. 

On  account  of  the  high  resistance  of  the  alcoholic  solutions  of  the 
acids,  it  was  found  necessary  to  make  use  entirely  of  the  cylindrical  type 
of  conductivity  cell.  The  method  of  obtaining  the  constants  of  these 
cells  was  described  by  White3  and  by  Wightman.4 

1Amer.  Chem.  Journ.,  44,  156  (1910);  46,  56  (1911);  48,  320,  411  (1912);  60,  1  (1913). 

2Journ.  Amer.  Chem.  Soo.,  36,  2247-2249  (1914). 

'Amer.  Chem.  Journ.,  42,  527  (1909).  4/Wd.,  42,  527  (1909) ;  44,  64  (1911). 


Conductivity  of  Organic  Acids  in  Ethyl  Alcohol.  131 

Since  the  percentage  temperature  coefficients  of  conductivity  for 
substances  dissolved  in  alcohol,  as  well  as  the  coefficient  of  expansion 
of  the  alcohol  itself,  are  so  large,  it  is  necessary  to  regulate  the  tempera- 
ture as  closely  as  possible.  This  was  done  by  the  combination  of  a 
specially  devised  gas-regulator  and  thermo-regulator.  These  have 
already  been  described  in  a  paper  by  Davis  and  Hughes.1 

The  constant-temperature  baths  were  of  the  improved  form  designed 
by  Davis,2  of  this  laboratory.  These  baths  are  of  about  60  liters 
capacity  and  are  made  of  copper.  Heat  is  applied  to  a  heavy  iron 
pipe  outside  the  bath,  and  water  is  kept  circulating  through  this  pipe 
by  means  of  propellers.  Only  a  small  portion  of  the  water  in  the  ther- 
mostat comes  into  immediate  contact  with  the  heated  surface,  and  this 
portion  is  subsequently  mixed  with  the  main  body  of  water,  thereby 
securing  much  more  even  distribution  of  temperature.  In  these  baths 
ordinarily  the  temperature  does  not  vary  more  than  0.02°,  which  is 
sufficiently  constant  for  our  purpose.  With  greater  precautions  as  to 
insulation  against  changes  in  temperature,  and  with  further  modifi- 
cation of  the  thermo-regulator,  the  variation  can  be  decreased  to  a  few 
thousandths  of  a  degree. 

A  TV-horsepower  direct-current  motor  served  as  the  source  of  power 
for  the  stirrers,  the  power  being  transmitted  by  belt  drives.  The  motor 
proved  to  be  a  great  improvement  over  the  hot-air  engines  formerly 
used. 

The  thermometers  were  of  the  differential  Beckman  type,  and  were 
carefully  compared  with  a  standard  Reichsanstalt  thermometer,  which 
had  been  calibrated  also  at  the  United  States  Bureau  of  Standards. 

The  resistance-box  which  was  used  throughout  this  entire  investi- 
gation had  also  been  calibrated  at  the  Bureau  of  Standards.  A  very 
fine  Kohlrausch  slide-wire  bridge  was  employed,  by  means  of  which  it 
was  possible  to  read  distances  on  the  slide-wire  corresponding  to  tenths 
of  a  millimeter  (the  total  length  of  the  wire  was  5  meters). 

Flasks,  pipettes,  and  burettes  for  measuring  purposes  were  in  all 
cases  carefully  calibrated. 

PROCEDURE. 

The  solutions  of  the  organic  acids  in  alcohol  were  made  up  in  200  c.c. 
Jena  measuring  flasks  calibrated  for  25°.  The  quantity  of  dried  and 
purified  acid  necessary  to  make  a  solution  of  the  required  normality 
was  weighed  out  on  a  watch-glass,  and  was  washed  off  through  a  funnel 
into  one  of  the  measuring  flasks.  The  flask  was  shaken  until  all  the 
acid  had  dissolved;  it  was  then  filled  to  the  neck  with  alcohol  and  sus- 
pended in  the  25°  bath.  When  temperature  equilibrium  was  reached 
it  was  filled  to  the  mark.  In  the  meantime  a  conductivity  cell  was 
thoroughly  washed  with  pure  alcohol  and  dried  with  filtered  dry  air. 

'Zeit.  phys.  Chem.,  85,  519  (1913).  'Carnegie  Inst.  Wash.  Pub.  No.  210,  121  (1915). 


132          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

It  was  then  nearly  filled  with  the  solution  which  had  been  made  up  and 
placed  in  the  15°  bath.  The  measuring-flask  containing  the  remainder 
of  the  solution  was  returned  to  the  25°  bath.  In  this  way  four  dilutions 
of  each  acid  were  prepared— N/8,  N/32,  N/128  and  N/512;  and  the 
four  cells  containing  them  placed  in  the  15°  bath.  In  addition,  a  fifth 
cell  was  employed  to  measure  the  conductivity  of  the  pure  solvent. 

After  the  cells  had  remained  in  the  15°  bath  for  at  least  an  hour,  the 
conductivities  of  the  solutions  were  measured.  Titrations  of  the  acids 
against  the  standard  ammonia  were  made  simultaneously  with  the 
conductivity  measurements.  For  this  purpose  10  c.c.  of  the  solution  in 
question  were  taken,  by  means  of  a  pipette,  from  the  proper  measuring 
flask  in  the  25°  bath.  A  second  titration  in  each  case  served  to  confirm 
the  results  of  the  first. 

When  the  conductivities  of  the  four  solutions  and  of  the  alcohol  had 
been  determined  at  15°,  the  cells  were  removed  to  the  25°  bath;  and 
after  the  lapse  of  an  hour  measurements  were  made  as  before.  The 
same  procedure  was  followed  for  the  determinations  at  35°. 

At  first  thought  it  would  seem  probable  that  keeping  the  solutions  in 
the  flasks  at  a  constant  temperature  (25°),  and  subjecting  the  solutions 
in  the  cells  to  changes  in  temperature  (15°  to  35°)  would  produce  a 
change  in  the  rates  of  esterification.  If  this  were  true,  the  normalities 
of  the  solutions  in  the  cells  would  be  different  from  the  values  obtained 
by  the  titration  of  the  solutions  in  the  flasks,  and  a  considerable  error 
would  be  introduced.  It  was  found,  however,  that  there  was  no  appre- 
ciable difference  in  the  amount  of  acid  present  at  any  moment  in  a  given 
solution,  whether  the  solution  was  kept  in  the  25°  bath  continuously 
for  8  hours,  or  whether  it  was  transferred  from  one  bath  to  another 
during  this  time.  The  reason  for  this  is  no  doubt  to  be  found  in  the 
extremely  slow  rate  at  which,  under  the  conditions  of  this  investigation, 
esterification  takes  place. 

Although  it  is  reasonably  certain  that  variation  in  the  temperature 
has  no  measurable  effect  upon  the  rate  of  esterification  in  alcoholic 
solutions  of  the  organic  acids,  this  variation  does  alter  to  a  considerable 
extent,  the  volume,  and  therefore  the  concentration,  of  these  solutions. 
For  example,  a  solution  which  has  a  volume  of  1,000  c.c.  at  25°,  con- 
tracts to  989.23  c.c.  when  cooled  to  15°,  and  expands  to  1,011.14  c.c. 
when  warmed  to  35°.  Because  of  this  fact  a  correction  had  to  be 
applied  to  the  volume  calculated  from  titration,  before  molecular  con- 
ductivity can  be  estimated  at  15°  and  35°.  This  correction  was  made 
in  the  following  way:  Let  us  suppose  that  the  normality  of  a  given 
solution  at  25°,  as  determined  by  titration  against  ammonia,  is  N2&. 
The  normality  at  15°,  Ni5,  would  then  be  expressed  by  the  ratio 

and  the  true  volume  at  15°  (i.  e.,  the  number  of  liters  which 


N25 


0.98923; 

contain  a  gram  molecular  weight  of  the  dissolved  acid)  would  be  the 


Conductivity  of  Organic  Acids  in  Ethyl  Alcohol.  133 

0  98923 
reciprocal  of  this  ratio,  or  --^  --     Similarly,  the  normality  at  35° 

NM  1.01114 

would  be  .  „..  -..  .  and  the  volume  would  be  —  ^  -- 

RESULTS. 

In  the  table  of  conductivity  results  (table  75),  Vm  is  the  volume  at 
which  the  solutions  were  made  up;  Vc  is  the  corrected  volume.  The 
corrections  were  applied  in  the  manner  just  described,  both  for  expan- 
sion or  contraction  of  the  alcohol  and  for  change  in  the  concentration 
of  the  acid  due  to  the  formation  of  ester.  Molecular  conductivity, 
Mt,  was  calculated  in  the  usual  manner. 

A  word  of  explanation  should  be  added  in  regard  to  the  method  of 
calculating  temperature  coefficients  of  conductivity.  The  formulae  gen- 
erally employed  in  the  determination  of  these  coefficients  are  T  = 

rr\ 

~  and  A  =  —  where  M*  and  Me  represent  the  molecular  conductivi- 


jt  —  ~ 

ties  of  the  same  solution  at  t  and  t',  respectively  (t'X);  T  is  the  tem- 
perature coefficient  expressed  in  conductivity  units,  and  A  is  the  per- 
centage temperature  coefficient.  But  these  formulae  as  such  are  not 
applicable  to  alcoholic  solutions  of  organic  acids,  because  the  true 
volume  (Ve)  of  a  given  solution  is  different  at  different  temperatures 
(due  to  the  causes  already  described).  For  example,  in  the  case  of 
benzoic  acid  (see  table  75),  the  N/8  solution  was  found  to  have  a 
volume  of  7.945  at  15°,  of  8.05  at  25°,  and  of  8.15  at  35°.  Before  the 
conductivities  of  these  three  solutions  can  be  compared  they  must  be 
reduced  to  values  corresponding  to  the  same  volume.  For  the  sake  of 
simplicity  Vc  at  25°  is  taken  as  the  standard  of  reference.  The  reduc- 
tion of  MC  at  15°  to  Vc  at  25°  is  made  in  the  following  way:  The  specific 

(0  0028Q2  \ 

7945     =  0.000364  J  is  multiplied  by  the  differ- 

ence in  volume  at  15°  and  25°  (8.05-7.945  =  0.105).  The  product 
(0.000038)  is  added  to  the  molecular  conductivity  at  15°  (0.002892), 
giving  0.002930,  which  represents  the  molecular  conductivity  at  15° 
of  a  solution  of  volume  8.05.  This  value  (0.002930)  and  the  value  of 
Mr  at  25°  (0.004073)  can  then  be  substituted  for  M<  and  Me  in  the  above 
equations,  and  T  and  A  are  found  to  be  0.0001143  and  3.90  respectively. 

(0  005444  \ 

'  s  JF  —  =  0.000668  J    is 

multiplied  by  the  difference  in  volume  at  25°  and  35°  (8.15-8.05  =  0.10). 
The  product  (0.000067)  is  subtracted  from  the  molecular  conductivity 
at  35°  (0.005444),  giving  0.005377,  which  represents  the  molecular 
conductivity  at  35°  of  a  solution  of  volume  8.05.  This  value  (0.005377) 
and  the  value  of  p,  at  25°  (0.004073)  can  then  be  substituted  for  Me 
and  M*  in  the  formulas,  and  T  and  A  are  determined  to  be  0.0001304  and 
3.20  respectively. 


134          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


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138 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


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3 


Conductivity  of  Organic  Adds  in  Ethyl  Alcohol. 


DISCUSSION  OF  THE  RESULTS. 


139 


The  most  striking  feature  of  the  conductivities  of  the  organic  acids 
in  alcohol,  as  shown  by  an  examination  of  table  75,  is  their  extremely 
small  value.  Wakeman,1  in  the  course  of  his  work  on  alcohol-water 
mixtures,  plotted  curves  of  conductivity  of  the  organic  acids  against 
percentage  alcohol,  and,  on  extending  the  curves  in  the  direction  of 
100  per  cent  alcohol,  found  that  they  apparently  approached  zero  con- 
ductivity as  a  limit.  As  can  be  seen  from  our  results,  the  conduc- 
tivities do  not  actually  approach  zero,  but  they  never  exceed  2,  and 
in  the  great  majority  of  cases  fall  below  unity.  On  account  of  the 
extremely  high  resistances  offered  by  alcoholic  solutions  of  the  organic 
acids,  Wildermann  abandoned  the  Kohlrausch  method  as  a  means  of 
studying  the  conductivity  of  these  solutions.  The  difficulty  which  he 
experienced  was  obviated  by  the  authors,  by  the  use  of  cells  with  much 
smaller  constants  than  those  he  employed.  Even  with  this  improve- 
ment, it  was  found  necessary  to  discard  all  of  the  very  weak  organic 
acids,  such  as  the  members  of  the  acetic-acid  series. 

The  conductivities  of  the  organic  acids  in  alcohol  are  several  hundred 
times  smaller  than  the  conductivities  of  the  same  acids  in  water.  When 
we  consider  the  fact  that  alcohol  has  from  one-fourth  to  one-fifth  the 
dissociating  power  of  water,  as  shown  by  the  dissociation  of  strong 
electrolytes  in  these  solvents,  the  above  fact  does  not  at  present  seem 
to  admit  of  any  very  satisfactory  explanation. 

The  percentage  temperature  coefficients  of  conductivity  vary  from 
2.5  to  4  per  cent,  and  decrease  with  rise  in  temperature.  They  are  uni- 
formly higher  than  the  corresponding  values  in  aqueous  solutions, 
which  range  from  1  to  2.5  per  cent.  In  the  light  of  the  work  that  has 
been  done  in  this  laboratory,  these  results  suggest  the  following 
possibilities : 

(1)  It  might  be  that  there  is  greater  alcoholation  in  alcoholic  solution 
than  hydration  in  aqueous  solution,  and  that  the  amount  of  alcohola- 
tion decreases  with  rise  in  temperature.     This  view,  however,  is  not 
substantiated  by  previous  experiences.     The  alcoholates  may  be  more 
unstable  with  rise  in  temperature  than  the  hydrates,  but  alcohol  seems 
in  general  to  have  far  less  power  to  combine  with  dissolved  substances 
than  water. 

(2)  If  the  temperature  coefficient  of  fluidity  of  alcohol  were  greater 
than  that  of  water,  the  ions  in  alcoholic  solutions  would  have  greater 
freedom  of  movement  than  in  aqueous  solutions.    This  would  account 
for  the  larger  temperature  coefficients  of  conductivity  in  alcoholic 
solutions.    Unfortunately,  such  is  not  the  case.     The  temperature 
coefficient  of  fluidity  of  water  is  greater  than  that  of  alcohol,  as  can 

'Zeit.  phys.  Chem.,11,  49  (1893) ;  Itid.,  14,  247  (1894). 


140         Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


be  seen  from  the  following  values  calculated  from  the  results  of  Wight- 
man,  Davis,  Holmes,  and  Jones:1 

TABLE  76. — Fluidities  and  temperature  coefficients  of  fluidity. 


Fluidities. 

Temperature  coef- 
ficients of  fluidity. 

15° 

25° 

35° 

15  to  25° 

25  to  35° 

Water  

88.18 
77.45 

112.23 
94.88 

138.89 
114.80 

2.73 
2.25 

2.37 
2.10 

Alcohol  .... 

BENZOIC  ACID 


MALEIC  ACID 


15° 


2.0 


a- 

'>  12 


"0.8 


Temperature. 
Fio.  13. 


35° 


Temperature 
Fio.  14. 


Nevertheless,  it  is  possible  that  the  greater  expansion  of  alcohol  with 
rise  in  temperature  affords  the  ions  greater  freedom  of  movement. 

(3)  A  suggestion  which  seems  to  be  far  more  probable  than  either 
of  the  preceding,  is  that  the  association  of  alcohol  is  affected  to  a  less 
extent  by  rise  in  temperature  than  is  the  association  of  water.  If  this 
be  true,  the  dissociating  power  of  alcohol  relative  to  that  of  water  will 

lJourn.  Chim.  Phys.,  12,  406  (1914). 


Conductivity  of  Organic  Acids  in  Ethyl  Alcohol. 


141 


become  greater  with  rise  in  temperature.  This  would  explain  the 
greater  temperature  coefficients  of  conductivity  of  alcoholic  solutions. 
The  increase  in  conductivity  with  rise  in  temperature  can  be  seen 
from  figures  13  and  14.  The  curves  have  very  much  the  appearance 
of  those  for  aqueous  solutions.  This  suggests  the  thought  that  perhaps 
the  increase  in  molecular  conductivity  in  alcohol  with  rise  hi  tempera- 
ture is,  as  in  aqueous  solutions,  a  parabolic  function,  and  that  the 
Euler  equation, 

£—  btz 


applies  to  both.  This  will  be  tested  in  later  work  by  determining  the 
conductivities  of  some  of  the  acids  at  temperatures  other  than  15°,  25°, 
and  35°,  and  comparing  the  results  obtained  with  those  calculated 
from  the  above  equation. 


BENZOIC  ACID 


MALEIC  ACID 


7 


Log  volume 
FIG.  15. 


i  z 

Log  volume 

Fio.  16. 


The  effect  of  increase  in  dilution  is  to  increase  the  molecular  con- 
ductivity. The  increase  in  conductivity  in  many  cases  is  almost  pro- 
portional to  the  volume.  This  relation  is  shown  graphically  in  figures 
15  and  16. 

A  knowledge  of  the  extent  to  which  organic  acids  are  dissociated  in 
alcoholic  solution  would  be  highly  desirable.  It  is  hoped  that  a  method 


142          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

for  the  determination  of  percentage  dissociation  will  be  worked  out 
here  in  the  near  future.  Goldschmidt1  obtained  values  for  the  limiting 
conductivities  of  several  organic  acids  in  alcohol  from  the  /JLX  values  of 
their  sodium  salts.  These  varied  from  83  to  93,  depending  on  the 
nature  of  the  acid.  It  has  not  yet  been  found  practicable  to  determine 
the  limiting  conductivities  of  the  organic  acids  studied  in  this  investi- 
gation, but  as  a  result  of  Goldschmidt's  work  it  is  certain  that  they  do 
not  differ  greatly  from  90.  If  this  be  the  case,  the  dissociation  of  the 
organic  acids  investigated  by  the  authors,  as  determined  by  conductivity, 
do  not  in  any  case  exceed  2  per  cent,  even  in  N/512  solutions. 

RELATION  BETWEEN  COMPOSITION  AND  CONDUCTIVITY. 

In  his  classical  study  of  the  conductivity  of  aqueous  solutions  of 
organic  acids,  to  which  reference  has  already  been  made,  Ostwald2 
pointed  out  a  number  of  relations  between  composition  and  molecular 
conductivity.  The  validity  of  these  relations  has  been  confirmed  by 
the  work  which  has  been  in  progress  in  this  laboratory  for  the  past 
fifteen  years.  In  view  of  this,  an  examination  of  the  results  obtained 
for  alcoholic  solutions,  in  the  attempt  to  discover  similar  relations, 
should  prove  to  be  interesting. 

Take  the  following  compounds: 

Maleic  acid.  Fumaric  acid.  Mesaconic  acid.  Itaconic  acid. 

H-C-COOH      HOOC-C-H      HOOC-C-CH,       CH2 

II  II  II  II 

H-C-COOH  H-C-COOH     H-C-COOH  C-COOH 

I 
CH2-COOH 

The  conductivity  of  maleic  acid  (table  75)  is  many  times  that  of 
fumaric.  This  fact  is  in  keeping  with  the  results  obtained  in  aque- 
ous solution,3  and  with  the  present  conception  of  the  configuration  of 
these  acids.  Mesaconic  acid  is  a  methyl  substitution  product  of 
fumaric  acid,  and  its  conductivity  is  of  the  same  order  of  magnitude 
as  that  of  fumaric  acid.  Itaconic  acid,  which  is  isomeric  with  mesaconic 
acid,  but  which  has  very  different  constitution,  shows  much  higher  con- 
ductivity. 

Malonic  acid,  at  a  volume  at  32  and  at  25°,  has  a  molecular  conduc- 
tivity of  0.055.  Under  the  same  conditions,  ethylmalonic  acid  has  a 
conductivity  of  0.083;  diethylmalonic  0.080;  propylmalonic  0.105; 
dipropylmalonic  0.090;  butylmalonic  0.036;  allylmalonic  0.039;  and 
benzylmalonic  0.062. 

The  above  results  show  that  the  introduction  of  an  ethyl  group 
increases  the  conductivity,  while  the  introduction  of  a  second  ethyl  group 

'Zeit.  Elektrochem.,  15,  4  (1909);  Zeit.  phya.  Chem.,  30,  627  (1910);  81,  30  (1912).     Zeit. 
Elektrochem.,  20,  473  (1914);  Zeit.  phys.  Chem.,  89,  129  (1914). 
!Zeit.  phys.  Chem.,  3,  170,  241,  369  (1889). 
'Carnegie  Inst.  Wash.  Pub.  No.  170,  113  (1912). 


Conductivity  of  Organic  Acids  in  Ethyl  Alcohol.  143 

tends  to  decrease  the  conductivity  of  the  ethylmalonic  acid.  Propylma- 
lonic  acid  has  uniformly  higher  conductivity  than  ethylmalonic,  and  the 
conductivity  of  the  dipropyl  acid  is  uniformly  higher  than  that  of  the 
diethyl.  Just  as  diethylmalonic  acid  has  smaller  conductivity  than 
ethylmalonic,  so  dipropylmalonic  acid  has  smaller  conductivity  than 
propylmalonic. 

Butylmalonic  and  allylmalonic  acids,  at  the  dilution  in  question, 
have  smaller  conductivities  than  malonic  acid  itself;  but  as  the  dilution 
increases,  the  conductivity  of  allylmalonic  acid  becomes  greater  than 
that  of  malonic  acid. 

Benzylmalonic  acid  has  greater  conductivity  than  malonic,  but 
less  than  ethylmalonic  acid.  This  is  especially  interesting,  in  consid- 
eration of  the  fact  that,  in  general,  a  phenyl  derivative  of  an  acid  has 
much  greater  conductivity  than  the  corresponding  methyl  derivative; 
e.  g.,  the  conductivity  of  acetic  acid  in  alcohol  is  so  small  that  it  can  not 
be  accurately  measured.  On  the  other  hand,  benzoic  acid  has  a  conduc- 
tivity of  0.014.  Again,  benzilic  acid  having  the  formula 

C6H6V     /OH 
C 


has  a  much  greater  conductivity  than  the  corresponding  methyl  deriva- 
tive, oxyisobutyric  acid 

CH3X      /OH 
C 


This  is  in  keeping  with  the  results  obtained  for  these  acids  in  aqueous 
solution,1  and  with  the  general  observation  that  the  replacement  of  a 
methyl  group  by  a  phenyl  group  increases  the  conductivity. 

Phenylpropiolic  acid,  C6H5  —  C  =  C  —  COOH,  has  a  conductivity  many 
times  larger  than  that  of  cinnamic  acid 

C6H5  -  CH  =  CH  -  COOH 

This  is  in  accord  with  what  was  found  for  these  same  acids  in  aqueous 
solutions. 

Take  the  following  series: 


Benzoic  acid  ............  CeH,COOH  1.3.5  dinitrobenzoic  acid.  .  . 

m-chlorbenzoic  acid  .....  C  J^CICOOH  o-toluic  acid 

m-nitrobenzoic  acid  .....  CeBUNOsCOOH  p-toluic  acid  ............  CeEUCHjCOOH 

The  conductivities  of  benzoic  acid  and  of  orthotoluic  acid  are  about 
equal,  whereas  the  conductivity  of  p-toluic  acid  is  very  much  less. 
That  the  same  relation  holds  also  for  other  substituent  groups  may 

'Carnegie  Inst.  Wash.  Pub.  170,  115,  132  (1912). 


144          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

be  seen  from  the  results  obtained  a  year  ago  by  Wightman,  Wiesel,  and 
Jones.1  For  example,  the  conductivities  of  o-chlorbenzoic  acid  and  of 
o-nitrobenzoic  acid  are  approximately  the  same  as  that  of  benzoic,  but 
the  conductivities  of  the  corresponding  para-acids  are  considerably  less. 

It  is  well  established  that  in  aqueous  solution  the  conductivity  of 
benzoic  acid  is  somewhat  increased  by  the  introduction  of  methyl, 
chlorine,  or  the  nitro  group,  in  the  para  position,  and  enormously 
increased  by  the  introduction  of  one  of  these  groups  in  the  ortho  posi- 
tion. In  contradistinction  to  this,  the  effect  of  these  groups  upon  the 
conductivity  in  alcoholic  solution  appears  to  be  negative.  Just  the 
opposite  is  true  if  the  chlorine  or  nitro  group  occupies  the  meta  position, 
as  can  be  seen  from  the  data.  In  these  cases  the  conductivity  of  ben- 
zoic acid  is  somewhat  increased. 

The  introduction  of  a  second  nitro  group  into  m-nitrobenzoic  acid 
still  further  increases  the  conductivity.  While  benzoic  acid,  C6H5COOH, 

,COOH 
has  a  conductivity  of  only  0.015,  phthalic  acid,  C6H4\  a 

XCOOH(o) 

dicarboxy  derivative,  has  a  conductivity  of  0.108,  which  is  seven  times 
as  great.  This  same  relation  holds  in  aqueous  solutions.2 

Of  all  the  acids  studied  in  this  investigation  with  one  exception, 
picric  acid,  C6H2(NO2)3OH,  is  the  strongest.  The  only  stronger  acid  in 
alcoholic  solutions  is  sulphosalicylic,  which  has  a  conductivity  ap- 
proaching that  of  hydrochloric  acid  in  this  solvent.  This  is  in  general 
analogous  to  what  was  found  when  these  compounds  were  dissolved  in 
water,3  sulphosalicylic  acid  in  water  having  almost  exactly  the  same 
conductivity  as  hydrochloric  acid  in  that  solvent. 

In  order  to  compare  the  conductivities  of  the  above-named  organic 
acids  in  alcohol  with  the  conductivities  of  these  same  compounds  in 
water,  reference  must  be  had  to  publication  of  the  Carnegie  Institution 
of  Washington  No.  170. 

'Journ.  Amer.  Chem.  Soc.,  36,  2251-2252  (1914).  'Ibid..  120,  121. 

"Carnegie  Inst.  Wash.  Pub.  No.  170,  116,  133  (1912). 


CHAPTER  VIII. 

CONDUCTIVITIES,  TEMPERATURE  COEFFICIENTS  OF  CONDUCTIVITY. 

AND  PERCENTAGE  DISSOCIATION  OF  SOME  RATHER 

UNUSUAL  SALTS  IN  AQUEOUS  SOLUTION. 


BY  CHARLES  WATKINS. 


This  investigation  is  a  contribution  to  the  study  of  the  conductivity, 
temperature  coefficients  of  conductivity,  and  dissociation  of  electrolytes 
in  aqueous  solutions.  It  is  a  continuation  of  the  work  begun  by  Jones 
and  West1  in  1905,  and  which  has  been  in  progress  continuously  since 
that  time.  After  studying  thirty-two  substances,  organic  and  inorganic, 
with  special  reference  to  the  effects  of  temperature  over  a  range  of  0° 
to  35°,  Jones  and  West  reached  the  following  conclusions : 

1.  A  large  increase  in  conductivity  due  to  greater  ionic  mobility 
accompanies  a  rise  in  temperature. 

2.  A  rise  in  temperature  brings  about  a  slight  decrease  in  dissocia- 
tion.    This  is  in  accord  with  the  law  of  Dutoit  and  Aston,  connecting 
the  association  of  a  solvent  with  its  dissociating  power,  and  is  in  har- 
mony with  the  results  of  Ramsey  and  Shields,  showing  the  relation 
between  the  dissociating  action  of  water  and  rise  in  temperature. 

3.  Temperature  coefficients  of  conductivity  expressed  in  percentage 
units  decrease  with  rise  in  temperature. 

4.  The  temperature  coefficients  of  conductivity  expressed  in  con- 
ductivity units,  in  the  case  of  salts,  increase  with  rise  in  temperature, 
while  in  the  case  of  acids  they  decrease. 

Jones  and  Jacobson2  next  investigated  35  compounds  over  the  same 
range  of  temperature  used  by  West,  and  found  that: 

1.  The  molecular  conductivity  of  electrolytes  in  aqueous  solutions 
increase  as  a  parabolic  function  of  the  temperature  from  0°  to  35°. 

2.  Hydrolysis  is  a  source  of  error  in  obtaining  the  value  of  nx. 

3.  The  conductivity  of  water  is  a  linear  function  of  the  temperature. 

4.  Salts  strongly  hydrated  in  solution  show  greater  increase  in  con- 
ductivity with  rise  in  temperature  than  salts  that  are  slightly  hydrated. 

Aqueous  solutions  of  organic  acids  were  the  subject  of  several  inves- 
tigations by  Jones  and  his  co-workers.  White,3  while  engaged  in  this 
work,  found  that  the  percentage  temperature  coefficients  of  conduc- 
tivity of  most  of  the  organic  acids  examined  were  small,  and  decreased 
with  rise  in  temperature  and  with  increase  in  dilution.  The  rate  of 
decrease  of  these  coefficients,  expressed  in  conductivity  units,  with  rise 
hi  temperature,  indicates  that  these  acids  are  much  less  hydrated  than 
mineral  acids.  In  the  case  of  the  amino  acids,  the  percentage  coeffi- 
cients of  conductivity  are  very  large.  Internal  salt  formation  was 

'Amer.  Chem.  Journ.,  34,  357  (1905).  *Ibid.,  40,  355  (1908).  >Ibid.,  42,  520  (1909). 

145 


146          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

suggested  as  the  cause  of  these  high  values.  He  also  shows  that  the 
conductivity  of  most  organic  acids  is  a  parabolic  function  of  the  tem- 
perature, but  concludes  that  no  general  statement  can  be  made  con- 
cerning the  change  in  dissociation  with  temperature  change,  since  the 
behavior  of  solutions  of  organic  acid  is  not  in  accord  with  the  Thompson- 
Nernst  hypothesis  connecting  dissociating  power  and  dielectric  constant. 

Clover1  continued  this  work,  but  over  a  wider  range  of  temperature. 
Some  measurements  were  made  at  a  temperature  of  80°.  West2  and 
Wightman3  also  did  work  in  this  field.  The  results  obtained  by  these 
investigators  were  in  accord  with  the  findings  of  their  predecessors. 

Alum  and  other  soluble  salts  were  studied  by  Hosford  and  Jones. 
They  found  that  the  conductivity  of  an  alum  is  less  than  the  conduc- 
tivity of  the  constituent  salts.  This  is  evidence  in  favor  of  a  theory 
advanced  by  Jones  regarding  the  existence  of  double  salts,  as  such,  in 
concentrated  solutions. 

While  studying  the  conductivity  of  a  number  of  electrolytes  in 
aqueous  solutions,  Winston  and  Jones4  found  apparent  exceptions  to 
the  earlier  work,  viz,  an  increase  in  percentage  dissociation  with  rise 
in  temperature,  and  large  temperature  coefficients  when  there  were  no 
indications  of  large  hydration.  This  led  Winston  to  advance  the  theory 
that  inductive  action  takes  place  through  the  solvent  between  charged 
ions  and  neutral  molecules,  giving  rise  to  complex  molecules  and  ions 
in  solution. 

Wightman,6  Springer,6  and  Smith7  extended  the  work  previously 
done  in  this  laboratory  on  organic  acids.  The  last-named  investigator 
found  that  Euler's  equation 

2—  bf? 


holds  in  every  case  studied,  and  calls  attention  to  the  fact  that  isomeric 
acids  do  not  behave  similarly  with  regard  to  change  in  dissociation. 
That  the  migration  velocity  of  anions  is  a  function  of  the  number  of 
atoms  constituting  the  anion  was  indicated  by  his  experimental  work. 

Howard8  and  Shaeffer9  have  each  published  papers  on  conductivity 
work  carried  out  under  the  direction  of  Jones. 

In  this  brief  review  of  the  conductivity  studies  made  in  this  labora- 
tory, attention  has  been  called  to  work  done  only  on  aqueous  solutions. 
Similar  work  has  been  carried  out  by  Jones  and  his  collaborators,  using 
glycerol,  ethyl  alcohol,  and  formamid  as  solvents.  A  considerable 
number  of  investigations  involving  various  mixtures  of  different  solvents 
have  been  carried  out.  For  a  full  account  of  this  work,  Publications 
80,  180,  and  210  of  the  Carnegie  Institution  of  Washington  should  be 
consulted. 

1Amer.  Chem.  Journ.,  43,  187  (1910).     4Ibid.,  46,  368  (1911).     'Zfcid.,  50,  1  (1913). 
'Ibid.,  44,  508  (1910).  <-Ibid.,  48,  320  (1912).     "Ibid.,  48,  500  (1912). 

'Ibid.,  46,  56  (1911).  'Ibid.,  48,  411  (1912).      '/Wd.,  49,  207  (1913). 


Conductivities  of  Some  Unusual  Salts.  147 

EXPERIMENTAL. 
APPARATUS. 

Bridge  and  Rheostat. — All  readingswere  made  on  an  improved  circular 
slide-wire  bridge  constructed  by  the  Leeds  and  Northrop  Co.,  of  Phila- 
delphia. From  a  bridge  of  this  type,  readings  can  be  made  to  a  fraction 
of  a  millimeter.  The  resistance-box  had  been  recently  standardized. 
The  plugs  of  this  piece  of  apparatus  were  kept  clean  by  frequently 
rubbing  them  with  a  piece  of  soft  chamois-skin  moistened  with  alcohol. 

Cells. — The  cells  used  were  of  the  same  type  described  and  sketched 
in  previous  articles1  on  this  subject.  Since  all  cells  used  had  been  in 
service  for  several  years,  no  trouble  was  experienced  from  the  leaching 
out  of  soluble  constituents  from  the  glass.  On  account  of  the  rather 
wide  range  of  conductivities  shown  by  the  salts  studied,  it  was  impos- 
sible, in  every  instance,  to  use  the  same  cell  for  the  same  dilution. 

Constant  Temperature  Baths. — With  the  exception  of  the  zero  bath, 
constant  temperature  was  maintained  by  the  application  of  a  principle 
developed  here  by  Morse.2  This  is  best  described  in  his  own  words: 

"If  all  the  water  or  air  in  a  bath  is  made  to  pass  rapidly  (1)  over  a  con- 
tinuously cooled  surface  which  is  capable  of  reducing  the  temperature  slightly 
below  that  which  it  is  desired  to  maintain,  then  (2)  over  a  heated  surface  which 
is  more  efficient  than  the  cooled  one,  but  under  the  control  of  a  thermostat, 
and  (3)  again  over  the  cooled  surface,  etc.,  it  should  be  practicable  to  maintain 
in  the  bath  any  temperature  for  which  the  thermostat  is  set,  and  the  con- 
stancy of  the  temperature  should  depend  only  on  the  sensitiveness  of  the 
thermostat  and  the  rate  of  flow  of  the  water  or  air.  The  principle  is  a  general 
one  and  provides  for  the  maintenance  of  any  temperature  between  zero  and 
the  boiling-point  of  water." 

The  particular  type  of  bath  and  thermostat  used  was  described  fully 
by  Davis  and  Putnam  in  Publication  No.  210  of  the  Carnegie  Institu- 
tion of  Washington.  In  making  a  zero  bath  of  the  type  described  by 
Jones  and  Jacobson,3  it  was  found  necessary  to  reduce  the  ice  to  a  fine 
state  of  division  with  an  ice-shaver  before  a  temperature  of  0°  could 
be  obtained. 

Containing  Vessels. — All  bottles  and  measuring  flasks  were  of  Jena 
glass.  After  recalibrating  the  flasks  by  the  weight  method,  the  flasks 
and  the  bottles  were  allowed  to  stand  partly  filled  with  a  solution  of 
chromic  acid  for  several  weeks.  During  this  time  they  were  frequently 
shaken  to  insure  thorough  cleansing. 

SOLUTIONS. 

Water. — The  water  used  in  the  preparation  of  the  solutions  was 
purified  by  the  method  of  Jones  and  MacKay4  as  modified  by  Schmidt.6 
In  no  case  was  the  specific  conductivity  greater  than  1.7X10"6. 

'Carnegie  Inst.  Wash.  Pub.  No.  170,  6  (1912).  'Ibid.,  19,  90  (1897). 

*1W.,  198,  (1914).  'Ibid-,  17,  83  (1895). 

'Amer.  Chem.  Journ.,  40,  361  (1908). 


148          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

Salts. — Kahlbaum's  or  Merck's  best  products  were  selected  in  all  cases. 
These  were  recrystallized  from  conductivity  water  before  being  used. 

Standardization. — On  account  of  the  nature  of  the  substances  studied, 
it  was  necessary  to  standardize  most  of  the  solutions  by  analytical 
methods.  These  methods  will  be  taken  up  briefly  in  connection  with 
the  discussion  of  the  individual  salts.  Standardization  by  analytical 
means  is  open  to  the  objection  that  solutions  must  stand  for  some  time 
before  the  conductivity  measurements  are  made.  This  vitiates  results 
in  the  case  of  salts  which  are  hydrolyzed  or  undergo  slight  decom- 
position in  solution.  In  the  first  instance  alkali  resulting  from  hydroly- 
sis will  attack  the  glass  container,  as  was  shown  by  the  etching  of  a 
Jena  glass  bottle  in  which  a  N/8  solution  of  trisodiumphosphate  had 
been  standing  for  several  weeks.  By  standardizing  as  rapidly  as  is  con- 
sistent with  accuracy,  and  by  keeping  the  solutions  in  a  cool,  dark  place, 
their  deleterious  effects  are  minimized.  From  the  solution  which  had 
been  analyzed,  the  others  were  made  by  diluting  with  conductivity 
water.  Care  was  taken  that  no  volume  less  than  35  c.c.  should  be 
involved  in  these  measurements. 

CELL  CONSTANTS. 

The  cell  constants  were  frequently  determined  throughout  this  work 
in  the  usual  way,  by  means  of  a  solution  of  pure  potassium  chloride. 
For  the  cells  with  electrodes  some  distance  apart,  a  N/50  solution  was 
used.  The  value  found  by  Kohlrausch,  129.7  (Siemens  units)  at  25° 
was  taken  for  the  molecular  conductivity  of  this  solution.  For  the  cells 
having  their  electrodes  closer  together,  a  N/500  solution  was  employed. 
The  molecular  conductivity  of  this  had  been  determined  in  a  cell  the 
constant  of  which  had  been  found  by  the  use  of  the  N/50  solution. 

PRECAUTIONS. 

Before  proceeding  to  the  experimental  data  and  methods  of  calcu- 
lating them,  it  may  be  well  to  call  attention  to  a  few  precautions  which, 
if  taken,  may  assist  those  engaged  in  work  requiring  apparatus  and 
methods  somewhat  similar  to  those  used  in  this  investigation. 

(a)  Jacobson1  has  already  called  attention  to  a  method  for  the  preven- 
tion of  the  collection  of  air-bubbles  on  the  electrodes  at  25°  and  above. 
He  recommends  heating  the  cell  and  contents  a  few  degrees  higher 
than  the  temperature  at  which  the  conductivity  is  to  be  measured,  and 
then  cooling  them  down  to  the  desired  temperature.  In  the  light  of 
Shaeffer's2  work,  it  is  seen  that  in  the  case  of  certain  salts,  notably 
chromium  compounds,  such  treatment  brings  about  a  change  in  con- 
ductivity which  persists  for  months.  Shaeffer's  observation  was  veri- 
fied in  the  course  of  the  present  investigation.  It  was  noticed  that  the 
conductivity  of  a  solution  of  ammonium  chromate  was  somewhat 

"Amer.  Chem.  Journ.,  40,  366  (1913).  'Ibid.,  49,  240  (1908). 


Conductivities  of  Some  Unusual  Salts.  149 

different  at  25°  after  standing  in  the  25°  bath,  from  the  conductivity 
of  the  same  solution  when  brought  to  a  temperature  of  25°  from  0°  to 
15°.  It  was  found  best  to  remove  the  bubbles  by  tilting  the  cells 
slightly  and  partly  removing  the  electrodes.  If  this  is  done  once  or 
twice  while  the  solutions  are  coming  to  temperature,  no  further  collec- 
tion of  bubbles  will  be  noticed. 

(6)  The  first  few  cubic  centimeters  of  conductivity  water  should  be 
discarded  from  the  siphon  used  in  drawing  the  water  from  its  container. 
The  water  remaining  below  the  stopcock  is,  to  some  extent,  exposed 
to  laboratory  atmosphere,  and  if  some  time  has  clasped  since  the 
removal  of  the  last  portion,  considerable  difference  in  conductivity 
will  be  noticed  if  the  above  precaution  is  neglected. 

(c)  All  connections  in  the  wiring  system  used  with  the  conductivity 
apparatus  should  be  soldered  where  this  will  not  injure  the  apparatus. 
When  not  in  use,  the  ends  of  the  wires  for  connecting  the  cells  with  the 
bridge  should  be  protected  by  small  glass  tubes. 

(d)  The  noise  of  the  motor  used  for  stirring  the  baths  was  found  to 
interfere  with  the  reading  of  the  bridge.     This  trouble  was  in  a  great 
measure  overcome,  by  mounting  the  motor  on  heavy  rubber  sheeting, 
and  providing  washers  of  the  same  material  for  all  bolts  used  in  holding 
this  machine  in  place. 

RESULTS. 

In  tables  77  to  94  all  conductivities  are  expressed  in  Siemens  units, 
and  are  molecular  conductivities  —  gram-molecular  weights  having 
been  used  in  preparing  the  solutions.  These  molecular  conductiv- 

Vet 
ities  GO  were  calculated  from  the  equation  fi,  =  K^r,  where  K  is  the 

cell  constant,  V  the  volume  concentration,  R  the  resistance  indicated 
on  the  rheostat,  (a)  and  (fc)  the  two  arms  of  the  bridge.  The  percen- 

tage dissociation  (a)  was  calculated  from  the  equation  a  =  -  —  ,  where 

Moo 

Hx  is  the  highest  value  of  /xc  obtained.  The  temperature  coefficients 
expressed  in  conductivity  units  were  calculated  from  the  formula 

-  coefficient 


£2—  1 

in  which  (/i,)^  represents  ^  at  the  higher  temperature  fe,  and  (jj,,)ti  at 
the  lower  temperature  (ti).  The  coefficients  expressed  as  percentages 
were  calculated  from  this  formula 


For  every  measurement  shown  in  these  tables  three  bridge  readings 
involving  different  values  for  R  were  made,  and  the  mean  of  these 
readings  was  taken  as  the  basis  of  calculation. 


150 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


SODIUM  BROMATE. 

This  salt  was  twice  recrystallized  and  then  was  dried  at  100°  for 
several  hours;  after  cooling  over  phosphorus  pentoxide,  the  required 
amount  was  weighed  and  dissolved.  (See  table  77.) 

TABLE  77. — Sodium  bromate. 


Molecular  conductivity. 

Percentage  dissocia- 
tion. 

Temperature  coefficients. 

V 

0° 

15° 

25° 

0° 

15° 

25° 

0  to  15° 

15  to  25° 

cond. 

cond. 

units. 

p.ct. 

units. 

p.ct. 

8 

42.77 

63.62 

78.84 

69.9 

74.4 

74.8 

1.38 

3.22 

1.52 

2.39 

16 

44.94 

67.03 

83.38 

73.5 

78.4 

79.1 

1.47 

3.27 

1.63 

2.43 

32 

47.49 

70.72 

87.95 

77.7 

82.7 

83.5 

1.54 

3.24 

1.72 

2.43 

128 

49.64 

74.39 

92.67 

81.2 

87.0 

88.0 

1.64 

3.30 

1.82 

2.45 

512 

52.69 

78.76 

96.53 

86.2 

92.1 

91.6 

1.73 

3.28 

1.77 

2.25 

1,024 

53.10 

79.59 

99.26 

86.9 

93.1 

94.2 

1.76 

3.31 

1.96 

2.47 

2,048 

58.27 

85.45 

105.1 

95.3 

100 

99.5 

1.81 

3.10 

1.97 

2.30 

4,096 

61.10 

83.34 

105.3 

100 

100 

1.48 

2.42 

2.19 

2.63 

SODIUM  SULPHOCYANATE. 

The  original  solution  of  this  salt  was  standardized  by  treating  a  meas- 
ured volume  with  a  slight  excess  of  an  acidified  solution  of  silver  nitrate. 
The  precipitate  of  silver  sulphocyanate  was  filtered  into  a  Gooch  cru- 
cible, washed  and  weighed.  Duplicate  determinations  were  made  as 
hi  the  case  of  all  solutions  standardized  by  analysis.  (See  table  78.) 

TABLE  78. — Sodium  sulphocyanate. 


Molecular  conductivity. 

Percentage  dissociation. 

Temperature  coefficients. 

V 

0° 

15° 

25° 

35° 

0° 

15° 

25° 

35° 

0  to  15° 

15  to  25° 

25  to  35° 

cond. 

cond. 

cond. 

units. 

p.ct. 

units. 

p.ct. 

units. 

p.ct 

4 

46.79 

69.00 

84.36 

102.5 

80.5 

79.8 

78.2 

78.4 

1.48 

3.16 

1.53 

2.22 

1.82 

2.15 

8 

49.68 

72.57 

96.36 

107.8 

85.5 

83.9 

89.4 

82.5 

1.52 

3.05 

2.38 

3.27 

1.15 

l.lfl 

16 

50.96 

75.86 

94.90 

114.2 

87.7 

87.7 

88.0 

87.4 

1.65 

3.23 

1.91 

2.51 

1.93 

2.03 

32 

52.90 

79.14 

98.81 

118.6 

91.0 

91.4 

91.6 

90.8 

1.74 

3.28 

1.97 

2.48 

1.98 

2.00 

128 

55.47 

83.19 

102.8 

124.3 

95.4 

96.1 

95.3 

95.1 

1.84 

3.31 

1.96 

2.35 

2.15 

2.09 

512 

57.77 

85.81 

106.7 

129.7 

99.4 

99.2 

99.0 

99.3 

1.86 

3.21 

2.09 

2.44 

2.29 

2.14 

1,024 

57.67 

86.50 

107.8 

130.9 

99.4 

100 

100 

100 

1.92 

3.32 

2.13 

2.46 

2.31 

2.14 

2,048 

58.10 

86.22 

107.8 

130.6 

100 

1.87 

3.21 

2.16 

2.50 

2.28 

2.11 

SODIUM  THIOSULPHATE. 

Iodine  which  had  been  resublimed  in  the  presence  of  potassium  iodide 
was  weighed  into  small  flasks  with  tightly  fitting  glass  stoppers.  About 
2  gm.  of  potassium  iodide  and  0.5  c.c.  of  water  had  been  previously 
weighed  into  these  flasks.  After  the  weight  of  the  added  iodine  was 
determined,  the  flasks  were  opened  in  Erlenmeyer  flasks  containing  20 
c.c.  of  water  and  potassium  iodide.  This  solution  was  titrated  with  the 
thiosulphate.  Starch  solution  was  used  as  an  indicator.  (See  table  79.) 


Conductivities  of  Some  Unusual  Salts. 
SODIUM  DITHIONATE. 


151 


After  recrystallizing  from  conductivity  water,  10  c.c.  of  the  sodium 
dithionate  solution  was  evaporated  to  dryness  in  a  weighed  platinum 
dish.  After  heating  to  dull  redness  for  some  time,  the  dish  and  contents 
were  cooled  over  calcium  chloride  and  weighed.  From  the  weight  of 
the  sodium  sulphate  the  strength  of  the  original  solution  of  sodium 
dithionate  was  calculated.  (See  table  80.) 

SODIUM  PYROPHOSPHATE. 

This  salt  was  dehydrated  and  weighed  after  it  had  been  recrystallized. 
(See  table  81.) 

TRISODIUM  PHOSPHATE. 

All  salts  of  othophosphoric  acid  were  treated  in  the  same  manner. 
After  making  a  solution  of  approximately  the  required  strength  by 
weighing  the  salt,  it  was  standardized  by  the  method  of  Schmitz.1 

TABLE  79. — Sodium  thiosulphate. 


Molecular  conductivity. 

Percentage  dissociation. 

Temperature  coefficients. 

V 

0° 

15° 

25° 

35° 

0° 

15° 

25° 

35° 

0  to  15° 

15  to  25° 

25  to  35° 

cand. 

cond. 

cond. 

units. 

p.ct. 

units. 

p.ct. 

units. 

p.  c 

4 

77  90 

115  3 

143  3 

172  6 

59  2 

59  1 

58  7 

59  4 

2  49 

3.07 

2.80 

2.43 

2.93 

?  n 

8 

82.52 

129.8 

160.8 

196.0 

62.8 

66.5 

65.9 

67.4 

3.15 

3.81 

3.10 

2.38 

3.52 

2.1 

16 

94  10 

143  4 

178  4 

214.8 

71.6 

73.5 

73.1 

73.9 

3.28 

3.48 

3.50 

2.44 

3.64 

?  0 

32 

103.0 

153.2 

193.0 

235.0 

78.4 

78.5 

79.1 

80.8 

3.34 

3.24 

3.98 

2.60 

4.19 

2.1 

128 

113.1 

172  A 

214.9 

259.9 

86.0 

88.4 

88.0 

89.4 

3.95 

3.49 

4.24 

2.46 

4.50 

2.0 

512 

126.0 

189.0 

234.0 

283.8 

95.8 

96.9 

95.9 

97.6 

4.19 

3.32 

4.50 

2.38 

4.98 

2.1 

1,024 

124.8 

191.9 

239.1 

284.1 

94.9 

98.4 

97.9 

97.7 

4.47 

3.58 

4.70 

2.45 

4.50 

1.8 

2,045 

128.1 

194.1 

241.1 

284.3 

97.5 

99.5 

98.8 

97.8 

4.40 

3.43 

4.69 

2.41 

4.32 

1.7 

4,096 

131.4 

195.0 

244.0 

290.5 

100 

100 

100 

100 

4.23 

3.22 

4.90 

2.50 

4.65 

1.9 

TABLE  80. — Sodium  dithionate. 


Molecular  conductivity. 

Percentage  dissociation. 

Temperature  coefficients. 

V 

0° 

15° 

25° 

35° 

0° 

15° 

25° 

35° 

0  to  15° 

15  to  25° 

25  to  35° 

cond. 

cond. 

cond. 

units. 

p.  ct. 

units. 

p.  ct. 

units. 

p.c 

8 

91.28 

135.9 

167.8 

202.6 

65.6 

65.2 

64.6 

64.4 

2.97 

3.00 

3.18 

2.34 

3.49 

2.C 

16 

99.71 

147.6 

183.9 

220.3 

71.6 

70.8 

70.8 

70.1 

3.19 

3.19 

3.63 

2.45 

3.64 

1.9 

32 

107  8 

161.7 

200.5 

241.6 

77.5 

77.6 

77.3 

76.8 

3.59 

3.33 

3.88 

2.40 

4.11 

2.0 

128 

121.4 

180.9 

225.0 

272.5 

87.2 

86.8 

86.7 

86.7 

3.96 

3.26 

4.49 

2.40 

4.75 

2.1 

512 

130.6 

195.6 

242.8 

293.9 

93.8 

93.9 

93.6 

93.5 

4.33 

3.31 

4.72 

2.41 

5.11 

2.1 

1,024 

137.1 

203.6 

255.0 

308.6 

98.6 

97.7 

98.3 

98.1 

4.43 

3.23 

5.14 

2.52 

5.36 

2.1 

2,048 

139.1 

208.4 

259.4 

314.3 

100 

100 

100 

100 

4.61 

3.31 

5.10 

2.44 

5.49 

2.1 

4,096 

139.1 

207.3 

258.2 

312.7 

4.54 

3.26 

5.09 

2.45 

5.45 

2.1 

'Zeit.  anal.  Chem.,  45,  512  (1906). 


152 


Conditctivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


This  differs  from  the  usual  method,  in  that  the  magnesium  ammonium 
phosphate  is  precipitated  in  a  hot  solution,  by  the  slow  addition  of 
ammonia  to  an  acid  solution  of  the  phosphate,  to  which  an  excess  of 
"magnesia  mixture"  had  been  previously  added.  A  very  pure, 
coarsely  crystalline  precipitate,  which  settles  readily,  is  obtained  by  this 
method.  From  the  weight  of  the  magnesium  pyrophosphate  yielded 
by  this  precipitate  on  ignition,  the  strength  of  the  solution  of  the  sodium 
salt  was  calculated.  (See  table  82.) 

TABLE  81. — Sodium  pyrophosphate. 


Molecular  conductivity. 

Percentage  dissociation. 

Temperature  coefficients. 

0° 

15° 

25° 

35° 

0° 

15° 

25° 

35° 

Oto  15° 

15  to  25° 

25  to  3 

cond. 

cond. 

cond. 

units. 

p.  ct. 

units. 

p.  ct. 

units. 

v 

16 

100.2 

154.8 

194.5 

235.3 

45.6 

44.9 

45.3 

44.9 

3.63 

3.62 

3.97 

2.56 

4.08 

i 

32 

118.3 

181.5 

227.5 

276.1 

53.8 

52.6 

53.0 

52.7 

4.20 

3.55 

4.60 

2.53 

4.86 

1 

128 

159.0 

242.7 

304.0 

370.5 

72.3 

70.4 

70.9 

70.8 

5.57 

3.49 

6.13 

2.52 

6.64 

2 

512 

197.0 

302.8 

384.5 

466.7 

89.6 

87.9 

89.7 

89.1 

7.51 

3.81 

8.17 

2.69 

8.22 

2 

1,024 

211.0 

327.0 

410.9 

508.9 

96.0 

94.9 

95.9 

97.2 

7.73 

3.66 

8.39 

2.56 

9.80 

2 

2,048 

216.1 

336.9 

421.8 

517.4 

98.3 

97.7 

98.4 

98.8 

8.05 

3.72 

8.49 

2.52 

9.50 

2 

4,096 

219.7 

344.5 

428.5 

523.4 

100 

100 

100 

100 

8.26 

3.75 

8.40 

2.43 

9.49 

2 

TABLE  82. — Trisodium  phosphate. 


Molecular  conductivity. 

Percentage  dissocia- 
tion. 

Temperature  coefficients. 

0° 

15° 

25° 

0° 

15° 

25° 

Oto  15° 

15  to  25° 

cond. 

cond. 

units. 

p.  ct. 

units. 

p.  ct. 

32 

120.5 

195.3 

254.2 

61.5 

63.6 

66.8 

4.98 

4.13 

5.89 

3.01 

128 

163.7 

261.7 

338.1 

83.6 

85.2 

88.8 

6.53 

3.98 

7.64 

2.91 

512 

185.2 

292.8 

361.5 

94.6 

95.3 

95.0 

7.17 

3.87 

6.86 

2.34 

1,024 

195.8 

307.0 

380.5 

100 

100 

100 

7.41 

3.79 

7.35 

2.39 

TABLE  83.— Sodium  dihydrogen  phosphate. 


Molecular  conductivity. 

Percentage  dissociation. 

Temperature  coefficients. 

0° 

15° 

25° 

35° 

0° 

15° 

25° 

35° 

Oto  15° 

15  to  25° 

25  to  » 

cond. 

cond. 

cond. 

units. 

p.  ct. 

units. 

p.  ct. 

units. 

P 

8 

31.02 

47.47 

59.79 

72.40 

72.2 

72.4 

73.2 

71.9 

1.09 

3.52 

1.23 

2.59 

1.26 

2 

16 

33.75 

51.62 

64.45 

79.55 

78.5 

78.7 

78.9 

79.0 

1.19 

3.52 

1.28 

2.48 

1.51 

2 

32 

36.16 

55.24 

69.25 

84.25 

84.2 

84.3 

84.8 

83.7 

1.27 

3.51 

1.40 

2.53 

1.50 

2 

128 

40.13 

60.98 

75.39 

93.67 

93.4 

93.0 

94.7 

93.1 

1.38 

3.43 

1.44 

2.36 

l.SO 

2 

512 

415  .  00 

65.52 

81.59 

99.90 

100 

100 

100 

99.3 

1.50 

3.46 

1.60 

2.45 

1.83 

2 

1,024 

42.07 

64.95 

81.17 

100.6 

100 

1.52 

3.55 

1.62 

2.48 

1.94 

2 

Conductivities  of  Some  Unusual  Salts. 


153 


SODIUM  DIHYDROGEN  PHOSPHATE. 

The  solution  of  this  salt  was  standardized  by  the  method  discussed 
in  connection  with  the  trisodium  phosphate.  (See  table  83.) 

SODIUM  TUNGSTATE. 

The  original  solution  of  sodium  tungstate  was  standardized  by  pre- 
cipitating the  tungstic  acid  as  mercurous  tungstate.  On  ignition,  this 
yields  the  trioxide  of  tungsten,  which  can  be  weighed.  (See  table  84.) 

SODIUM  FORMATE. 

In  a  cold  acid  solution,  permanganate  acts  only  slowly  with  formic 
acid,  while  in  a  hot  solution  the  latter  is  lost  by  volatilization.  In  spite 
of  statements  to  the  contrary,  oxidation  by  means  of  permanganate  in 
an  alkaline  solution  is  a  most  unsatisfactory  process.  This  led  to  the 
use  of  an  indirect  method  of  standardization  for  the  sodium  formate 
solution.  An  excess  of  standard  permanganate  solution  was  measured 
from  a  burette  into  an  alkaline  solution  of  the  formate.  After  standing 

TABLE  84. — Sodium  tungstate. 


Molecular  conductivity. 

Percentage  dissociation. 

Temperature  coefficients. 

0° 

15" 

25° 

35° 

0° 

15° 

25° 

35° 

0  to  15° 

15  to  25° 

25  to  35° 

cond. 

cond. 

cond. 

units. 

p.ct. 

units. 

p.  ct. 

units. 

p.ct 

4 

59.56 

100.2 

54.8 

59.9 

2.71 

4.55 

8 

69.75 

112.7 

141.6 

172.0 

64.2 

67.3 

66.9 

67.3 

2.86 

4.10 

2.88 

2.56 

3.24 

2.2* 

16 

82.65 

126.7 

159.0 

193.2 

76.1 

75.7 

75.2 

75.7 

2.93 

3.54 

3.23 

2.54 

3.42 

2.  IE 

32 

87.53 

134.3 

168.7 

205.5 

80.6 

80.3 

79.8 

80.5 

3.12 

3.56 

3.44 

2.56 

3.68 

2.  IS 

128 

98.54 

151.5 

190.7 

233.5 

90.7 

90.5 

90.2 

91.4 

3.52 

3.57 

3.92 

2.58 

4.28 

2.24 

512 

107.25 

164.6 

207.4 

254.1 

98.7 

98.3 

98.1 

99.5 

3.82 

3.56 

4.28 

2.60 

4.67 

2.2fi 

1,024 

108.67 

167.3 

211.4 

255.2 

100 

100 

100 

100 

3.90 

3.56 

4.41 

2.63 

4.38 

2.07 

TABLE  85. — Sodium  formate. 


y 

Molecular  conductivity. 

Percentage  dissociation. 

Temperature  coefficients. 

0" 

15° 

25° 

35° 

0° 

15° 

25° 

35° 

0  to  15° 

15  to  25° 

25  to  35° 

1 

'• 

cond. 

cond. 

cond. 

units. 

p.ct. 

units. 

p.  ct. 

units. 

p.  ct 

4 

56.09 

84.34 

105.0 

126.4 

70.3 

71.0 

70.9 

68.6 

1.88 

3.35 

2.07 

2.45 

2.14 

2.03 

8 

62.13 

91.91 

113.6 

140.1 

77.9 

77.4 

76.7 

76.0 

1.84 

2.96 

2.16 

2.35 

2.65 

2.33 

16 

64.75 

97.82 

121.4 

147.4 

81.2 

82.4 

81.9 

80.0 

1.53 

2.36 

2.35 

2.41 

2.60 

2.14 

32 

67.72 

99.10 

130.2 

157.1 

84.9 

83.5 

87.9 

85.3 

2.09 

3.08     3.11 

3.13 

2.69 

2.00 

128 

72.48 

109.8 

137.7 

166.3 

90.9 

92.5 

92.9     90.3 

2.48 

3.42  j  2.80 

2.55 

2.86 

2.07 

512 

74.63 

109.6 

139.1 

166.7 

93.6 

92.4 

93.9 

90.5 

2.35 

3.14 

2.95 

2.69 

2.76 

1.99 

1,024 

75.65 

112.8 

140.9 

181.0 

94.8 

95.0 

95.1 

98.3 

2.47 

3.26 

2.81 

2.49 

4.01 

2.84 

2,048 

78.68 

114.8 

147.9 

184.7 

98.7 

96.7 

99.8  ilOO 

2.41 

3.06 

3.31 

2.88 

3.68 

2.49 

4,096 

79.73 

118.7 

148.1 

184.1 

100 

100 

100 

2.59 

3.24 

2.94 

2.49 

3.60 

2.43 

154 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


for  some  time,  a  known  weight  of  pure  sodium  oxalate  was  added.  The 
solution  was  then  acidified  and  heated.  A  standard  solution  of  per- 
manganate was  used  to  titrate  this  in  the  usual  way.  The  total  amount 
of  permanganate  less  that  which  is  equivalent  to  the  sodium  oxalate  is 
the  permanganate  used  for  the  oxidation  of  the  sodium  formate.  (See 
table  85.) 

SODIUM  CHROMATE. 

The  original  solution  of  this  salt  was  standardized  by  precipitating 
the  chromium  as  mercurous  chromate,  by  means  of  a  solution  of  mer- 
curous  nitrate.  On  ignition,  the  mercurous  chromate  leaves  a  residue 
of  chromic  oxide  which  can  be  weighed.  (See  table  86.) 

SODIUM  DICHROMATE. 

The  same  method  of  standardization  was  used  in  the  case  of  sodium 
dichromate  as  was  employed  with  the  chromate.  (See  table  87.) 

POTASSIUM  FERRICYANIDE. 

This  salt  was  dried  for  more  than  a  month  over  phosphorus  pentoxide. 
The  solution  was  prepared  by  weighing  the  required  amount  of  the 
dry  salt,  and  dissolving  in  the  usual  manner.  (See  table  88.) 

TABLE  86. — Sodium  chromate. 


V 

Molecular  conductivity. 

Percentage  dissociation. 

Temperature  coefficients. 

0° 

15° 

25° 

35° 

0° 

15° 

25° 

35° 

0  to    15° 

15  to  25° 

25  to  35° 

4 
8 
16 
32 
128 
612 
1,024 

74.76 
83.56 
90.74 
98.16 
110.40 
118.32 
114.60 

113.4 
125.3 
137.8 
148.6 
168.4 
176.2 
173.6 

63.2 

64.3 

cond. 
units. 
2.57 

p.  ct. 
3.43 

cond. 
units. 

p.  ct. 

cond. 
unite. 

p.  ct. 

156.5 
171.3 
185.1 
206.7 
219.5 
216.9 

187.9 
207.6 
224.3 
252.2 
270.1 
262.3 

70.6 
76.7 
82.9 
93.3 
100 

71.1 

78.2 
84.3 
95.5 
100 

71.3 

78.0 
84.3 
94.2 
100 

69.6 
76.8 
83.0 
93.4 
100 

2.78 
3.13 
3.36 
3.86 
3.85 
3.92 

3.32 
3.44 
3.42 
3.49 
3.25 
3.42 

3.12 
3.35 
3.65 
3.83 
4.33 
4.33 

2.48 
2.43 
2.45 
2.27 
2.45 
2.49 

3.14 
3.63 
3.92 
4.65 
5.06 
4.54 

2.00 
2.11 
2.11 
2.20 
2.30 
2.09 

TABLE  87. — Sodium  dichromate. 


V 

Molecular  conductivity. 

Percentage  dissociation. 

Temperature  coefficients. 

0° 

15° 

25° 

35° 

0° 

15° 

25° 

35° 

0  to  15° 

15  to  25° 

25  to  35° 

8 
16 
32 
128 
612 

92.18 
96.57 
101.25 
106.8 
106.1 

136.5 
143.5 
148.4 
158.3 
158.3 

87  2 

86  22 

cond. 
units. 
2  95 

p.  ct. 
3  20 

cond. 
units. 

p.  ct. 

cond. 
units. 

p.ct. 

176.3 
182.7 
194.7 
194.4 

211.4 
219.5 
233.5 
234.6 

91.3 
95.7 
100 

90.65 
93.74 
100 

90.71 
93.98 
100 

89.88 
93.32 
100 

3.13 
3.14 
3.43 
3.50 

3.24 
3.10 
3.21 
3.31 

3.28 
3.43 
3.64 
3.61 

2.28 
2.31 
2.29 
2.28 

3.51 

3.68 
3.88 
4.02 

1.99 
2.01 
1.99 
2.06 

Conductivities  of  Some  Unusual  Salts. 


155 


AMMONIUM  IODATE. 

A  very  dilute  solution  of  the  iodate,  while  cold,  was  carefully  treated 
with  sulphurous  acid.  When  an  amount  sufficient  to  remove  the 
brown  color  due  to  free  iodine  had  been  added,  the  solution  was  warmed 
until  the  odor  of  sulphur  dioxide  could  not  be  detected.  The  iodine 
was  then  determined  as  silver  iodide.  (See  table  89.) 

TABLE  88. — Potassium  ferricyanide. 


Molecular  conductivity. 

Temperature  coefficients. 

0° 

15° 

25° 

35° 

0  to  15° 

15  to  25° 

25  to  35° 

cond. 

cond. 

cond. 

units. 

p.  ct. 

units. 

p.  ct. 

units. 

p.  ct. 

8 

158.6 

230.2 

282.1 

4.77 

3.00 

5.19 

2.25 

16 

168.8 

247.3 

303.3 

360.1 

5.23 

3.09 

5.60 

2.26 

5.68 

1.87 

32 

181.4 

266.4 

326.4 

389.3 

5.66 

3.11 

6.00 

2.25 

6.29 

1.92 

128 

207.4 

308.2 

380.1 

451.7 

6.71 

3.23 

7.19 

2.33 

7.16 

1.88 

512 

229.2 

331.8 

410.1 

494.5 

6.83 

2.98 

7.87 

2.37 

8.40 

2.04 

1,024 

238.1 

354.5 

438.9 

529.9 

7.75 

3.25 

8.44 

2.37 

9.10 

2.07 

2,048 

244.4 

363.0 

449.5 

543.3 

7.90 

3.23 

8.65 

2.38 

9.38 

2.08 

4,096 

254.5 

375.3 

465.8 

564.3 

8.05 

3.16 

9.04 

2.40 

9.86 

2.11 

TABLE  89. — Ammonium  iodate. 


Molecular  conductivity. 

Percentage  dissociation. 

Temperature  coefficients. 

0° 

15° 

25° 

35° 

0° 

15° 

25° 

35° 

0  to  15° 

15  to  25° 

25  to  35° 

cond. 

cond. 

cond. 

units. 

p.  ct. 

units. 

p.  ct. 

unity. 

p.  ct. 

16 

48.17 

72.21 

89.97 

108.5 

85.1 

84.88 

84.7 

84.4 

1.60 

3.32 

1.77 

2.45 

1.85 

2.05 

32 

50.51 

76.35 

94.93 

114.6 

89.2 

89.74 

89.4 

89.1 

1.72 

3.40 

1.85 

2.43 

1.97 

2.07 

128 

54.34 

81.32 

101.1 

122.2 

96.0 

95.59 

95.2 

95.0 

1.79 

3.29 

1.97 

2.43 

2.11 

2.08 

512 

56.54 

85.07 

106.2 

128.6 

99.9 

100 

100 

100 

1.90 

3.35 

2.12 

2.49 

2.24 

2.10 

1,024 

56  60 

84  93 

105  9 

128  0 

100 

1  88 

3  32 

2  10 

2  47 

2  21 

2  08 

2,048 

55.99 

83.92 

104.6 

127.2 

1.86 

3.32 

2.07 

2.46 

2.26 

2.16 

TABLE  90. — Ammonium  dihydrogen  phosphate. 


Molecular  conductivity. 

Percentage  dissociation. 

Temperature  coefficients. 

0° 

15° 

25° 

35° 

0° 

15° 

25° 

35° 

0  to  15° 

15  to  25° 

25  to  35° 

cond. 

cond. 

cond. 

units. 

p.  ct. 

units. 

p.  ct. 

units. 

p.  ct. 

4 

37.56 

56.27 

69.66 

83.53 

66.4 

67.0 

66.5 

65.9 

.24 

3.30 

1.34 

2.37 

1.38 

1.98 

8 

41.18 

61.73 

76.50 

92.71 

72.8 

73.5 

73.1 

73.1 

.36 

3.30 

1.47 

2.39 

1.62 

2.11 

16 

44.45 

66.55 

82.57 

99.90 

78.6 

79.3 

78.9 

78.8 

.47 

3.30 

1.60 

2.40 

1.73 

2.09 

32 

46.72 

70.65 

87.73 

106.48 

82.6 

84.1 

83.8 

84.0 

.59 

3.40 

1.70 

2.41 

1.87 

2.13 

128 

50.74 

76.21 

94.89 

114.75 

89.8 

90.8 

90.7 

90.5 

.69 

3.33 

1.86 

2.45 

1.98 

2.09 

512 

53.39 

81.01 

99.73 

121.2 

94.5 

96.5 

95.3 

95.6 

.83 

3.42 

1.87 

2.31 

2.15 

2.15 

1,024 

56.50 

83.92 

104.6 

126.7 

100 

100 

100 

100 

.83 

3.23 

2.07 

2.46 

2.21 

2.11 

156 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


AMMONIUM  DIHYDROGEN  PHOSPHATE. 

This  solution  was  standardized  in  the  same  way  as  described  in 
connection  with  sodium  phosphate.     (See  table  90.) 

AMMONIUM  CHROMATE. 

The  method  used  for  standardization  was  identical  with  that  employed 
for  standardizing  the  chromate  of  sodium.     (See  table  91.) 

TABLE  91. — Ammonium  chromate. 


V 

Molecular  conductivity. 

Percentage  dissociation. 

Temperature  coefficients. 

0° 

15° 

25° 

35° 

0° 

15° 

25° 

35° 

0  to  15° 

15  to  25° 

25  to  35° 

4 
8 
16 
32 
128 
512 
1,024 

93.63 
102.5 
110.5 
119.1 
135.1 
144.6 
143.0 

137.0 
149.8 
163.1 
176.3 
201.2 
213.2 
212.6 

64  7 

64  2 

cond. 
unite. 
2  89 

p.ct. 
3  08 

cond. 
units. 

p.ct. 

cond. 
units. 

p.ct. 

181.9 
199.5 
215.4 
247  1 

215.4 
236.7 
257.2 
292  7 

70.8 
76.4 
82.3 
93  4 

70.3 
76.5 
82.7 
94  3 

69.3 
76.0 
82.1 
94  2 

68.8 
75.7 
82.2 
03  fi 

3.15 
3.50 
3.81 
4  40 

3.07 
3.16 
3.19 
3  25 

3.21 
3.64 
3.91 
4  59 

2.14 
2.23 
2.21 

2  28 

3.35 
3.72 
4.18 
4  56 

1.88 
1.89 
1.94 
1.84 

262.2 
261.7 

312.7 
312.3 

100 

100 

100        100 

4.57 
4.63 

3.16 
3.23 

4.90 
4.91 

2.29 
2.30 

4.95 
5.06 

1.88 
1.93 

TABLE  92. — Ammonium  sulphocyanate. 


Molecular  conductivity. 

Percentage  dissociation. 

Temperature  coefficients. 

0° 

15° 

25° 

35° 

0° 

15° 

25° 

35° 

0  to  15° 

15  to  25° 

25  to  35° 

cond. 

cond. 

cond. 

units. 

p.  ct. 

units. 

p.ct. 

units. 

p.ct. 

4 

60.46 

86.53 

106.0 

125.7 

83.9 

82.2 

82.0 

81.2 

1.73 

2.86 

2.05 

2.37 

1.97 

1.89 

8 

61.47 

89.84 

110.3 

131.7 

85.3 

85.3 

85.3 

85.0 

1.89 

3.07 

2.04 

2.27 

2.14 

1.94 

16 

63.58 

93.60 

114.4 

137.5 

88.3 

88.9 

88.5 

88.8 

2.00 

3.14 

2.08 

2.22 

2.31 

2.01 

32 

66.30 

97.48 

119.5 

143.4 

92.0 

92.6 

92.5 

92.6 

2.07 

3.12 

2.20 

2.26 

2.39 

2.00 

128 

68.58 

101.2 

124.2 

148.9 

95.2 

96.2 

96.1 

96.2 

2.17 

3.16 

2.29 

2.26 

2.47 

1.99 

512 

71.74 

105.2 

129.0 

154.5 

99.6   100 

99.8 

99.8 

2.23 

3.10 

2.48 

2.35 

2.55 

1.97 

1,024 

72.00 

105.2 

129.2 

154.8 

100 

100 

100 

2.21 

3.06 

2.40 

2.28 

2.56 

1.97 

TABLE  93. — Lithium  chromale. 


Molecular  conductivity. 

Percentage  dissociation. 

Temperature  coefficients. 

0° 

15° 

25° 

35° 

0° 

15° 

25° 

35° 

Oto  15° 

15  to  25° 

25  to  35° 

1 

cond. 

cond. 

cond. 

units. 

p.  ct.  ]  units. 

p.ct. 

units. 

p.  ct. 

8 

74.62 

112.3 

139.9 

169.1 

65.4 

67.3 

66.9 

67.2 

2.51 

3.36 

2.76 

2.45 

2.92 

2.08 

16 

82.62 

124.1 

154.3 

187.3 

72.4 

74.3 

73.8 

74.5 

2.76 

3.34 

3.02 

2.45 

3.30 

2.13 

32 

89.83 

136.6 

169.9 

205.5 

78.7 

81.8 

81.3 

81.7 

3.11 

3.46 

3.33 

2.43 

3.56 

2.09 

128 

101.6 

155.1 

193.0 

236.5 

89.0 

92.9 

92.3 

94.0 

3.56 

3.50 

3.79 

2.44 

3.34 

1.73 

512 

106.4 

164.5 

205.3 

249.0 

93.2 

98.5 

98.2 

99.0 

3.87 

3.60 

4.08 

2.48 

4.37 

2.12 

1,024 

108.1 

166.0 

206.2 

250.3 

94.7 

99.4 

98.6 

99.5 

3.85 

3.59 

4.02 

2.42 

4.41 

2.13 

2,048 

114.1 

168.3 

209.0 

251.4 

100 

100 

100 

100 

3.61 

3.16 

4.07 

2.41 

4.24 

2.02 

Conductivities  of  Some  Unusual  Salts. 
AMMONIUM  SULPHOCYANATE. 


157 


The  amount  of  sulphocyanate  present  in  the  original  solution  was 
found  by  weighing  the  silver  sulphocyanate  formed  on  treating  a  portion 
of  the  solution  with  silver  nitrate.  (See  table  92.) 

LITHIUM  CHROMATE. 

The  chromate  in  this  solution  was  determined  as  it  was  in  the  case  of 
sodium  chromate.  (See  table  93.) 

RUBIDIUM  IODIDE. 

The  iodine  was  determined  as  silver  iodide.     (See  table  94.) 

TABLE  94. — Rubidium  iodide. 


Molecular  conductivity. 

Percentage  dissociation. 

Temperature  coefficients. 

V 

0° 

15° 

25° 

35° 

0° 

15° 

25° 

35° 

0  to  15° 

15  to  25° 

25  to  35° 

cond. 

cond. 

cond. 

units. 

p.  ct. 

units. 

p.  Ct. 

units. 

p.  c(. 

4 

68.40 

97.08 

85.41 

82.9 

1.91 

2.79 

8 

70.48 

101.4 

122.6 

145.6 

88.0 

86.6 

85.6 

85.2 

2.06 

2.92 

2.01 

1.98 

2.30 

.87 

16 

72.59 

104.3 

127.2 

150.6 

90.6 

89.0 

88.8 

88.1 

2.11 

2.90 

2.29 

2.19 

2.34 

.83 

32 

74.75 

108.6 

131.7 

157.8 

93.3 

92.7 

92.0 

92.3 

2.25 

3.01 

2.31 

2.12 

2.61 

.98 

128 

78.51 

114.2 

139.7 

166.6 

98.0 

97.5 

97.6 

97.5 

2.38 

3.03 

3.57 

3.12 

2.69 

.92 

512 

79.92 

116.0 

142.0 

169.5 

99.8 

99.0 

99.2 

99.2 

2.40 

3.00 

3.60 

3.11 

2.75 

.93 

1,024 

80.08 

117.1 

143.1 

170.8 

100 

100 

100 

100 

2.46 

3.07 

3.70 

3.16 

2.77 

.93 

DISCUSSION  OF  RESULTS. 
CONDUCTIVITIES. 

Electrical  conductivity  in  solutions  of  electrolytes  depends  on  the 
number  of  ions  present  and  on  the  velocities  of  these  ions.  The 
velocities,  other  conditions  being  the  same,  depend  upon  the  size  and 
mass  of  the  ion  and  upon  the  viscosity  of  the  medium.  This  leads 
to  the  conclusion  that  a  salt  showing  a  high  conductivity  must  be 
dissociated  into  a  great  number  of  ions,  or  the  ions  in  solutions  must  be 
of  such  a  size  and  mass  that  they  have  a  great  velocity. 

It  can  be  seen  from  the  above  tables  that  certain  salts,  notably 
trisodium  phosphate,  sodium  pyrophosphate,  ammonium  chromate, 
and  potassium  ferricyanide,  show  very  high  conductivity.  The  first 
of  these  compounds  is  strongly  hydrolyzed  even  at  low  temperatures. 
The  breaking  down  of  complexes  by  hydrolysis  gives  rise  to  a  great 
number  of  ions  also  in  the  case  of  the  pyrophosphate.  It  is  interesting 
to  compare  the  conductivity  of  the  trisodium  phosphate  with  that  of 
the  corresponding  potassium  compound.  While  both  show  very  high 
conductivities  at  all  temperatures,  the  conductivity  of  the  potassium 
salt  is  greater  than  that  of  the  sodium.  We  would  expect  this  from 


158          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

the  fact  that  the  sodium  salt  crystallizes  with  12  molecules  of  water, 
indicating  great  hydration  in  solution;  while  the  potassium  salt  has  no 
water  of  crystallization,  which  indicates  only  slightly  hydrated  ions 
in  solution.  The  ammonium  chromate  is  somewhat  similar  to  the 
unhydrated  potassium  salt  mentioned  above,  in  that  it  carries  no  water 
of  crystallization,  and  would  therefore  be  expected  to  show  greater 
conductivity  than  the  corresponding  hydrated  salts  of  sodium  and 
lithium.  The  high  values  of  /j.  for  potassium  ferricyanide  is  largely 
due  to  the  great  number  of  ions  yielded  by  this  compound.  The  work 
of  Getman  and  Bassett1  indicates  the  production  of  6  ions  in  solutions 
of  this  salt.  By  comparing  the  conductivities  of  the  chromates  of 
ammonia  and  the  alkali  metals,  it  is  found  that  they  stand  in  the  fol- 
lowing order: 

Potassium  chromate  >  ammonium  chromate  >  sodium  chromate  > 
lithium  chromate. 

Ammonium  compounds,  as  a  rule,  show  higher  conductivity  than 
the  corresponding  potassium  salts.  This  does  not  seem  to  be  true  in 
the  case  of  the  chromates.  The  smaller  conductivities  of  lithium 
compounds  when  compared  with  compounds  of  sodium,  is  usually 
attributed  to  the  greater  hydration  of  the  lithium  ion  in  solution,  as 
indicated  by  the  greater  tendency  of  lithium  salts  to  crystallize  with 
water.1  The  chromate  of  lithium,  however,  crystallizes  with  1  mole- 
cule of  water,  while  the  chromate  of  sodium  contains  10  molecules. 
The  work  of  Jones  and  Bassett2  has  shown  that  many  substances  have 
greater  hydrating  power  than  is  indicated  by  the  water  of  crystalli- 
zation contained  in  them.  Suchmaybethe  case  with  lithium  chromate. 

Rubidium  iodide  shows  a  higher  conductivity  than  the  iodides  of 
the  other  alkali  metals.  Knowing  that  rubidium  has  a  greater  atomic 
volume  than  sodium  or  potassium,  we  might  expect  the  conductivity  to 
be  lowered  by  a  decrease  in  the  velocity  of  the  ions,  due  to  their  volume 
and  mass.  It  should,  however,  be  remembered  that  the  hydrating 
power  of  these  compounds  of  the  alkali  metals  decreases  with  increasing 
atomic  volume.  Just  as  sodium  salts  are  less  hydrated  than  lithium,  so 
rubidium  compounds  would  be  expected  to  hydrate  less  than  potassium. 
Thus,  an  apparent  exception  is  explained  by  the  theory  of  hydration. 

DISSOCIATIONS. 

As  a  means  of  determining  the  dissociation  of  salts  in  solution,  the 
conductivity  method  is  of  great  service,  but  it  is  far  from  perfect. 
Hydrolysis,  hydration,  and  polymerization  all  militate  against  obtaining 
a  true  value  for  /j.x.  Since  most  salts  show  one  or  more  of  the  above- 
named  phenomena,  it  is  certain  that  dissociations  calculated  from  con- 
ductivity data  are  in  most  cases  simply  close  approximations. 

'Carnegie  Inst.  Wash.  Pub.  No.  60,  p.  46  (1907).  *Amer.  Chem.  Journ.,  33,  562  (1905). 


Conductivities  of  Some  Unusual  Salts.  159 

As  a  rule,  salts  in  aqueous  solutions  are  more  dissociated  at  low 
than  at  high  temperatures.  This  is  in  accord  with  the  Thompson- 
Nernst  hypothesis  to  which  reference  has  already  been  made.  Some 
exceptions  to  this  rule  have  been  found  by  other  investigators.  Barium 
nitrate,  cadmium  iodide,  lead  nitrate,  and  uranyl  acetate  have  been 
found  to  show  an  increase  in  dissociation  at  higher  temperatures. 
Shaeffer1  has  called  special  attention  to  the  anomalous  behavior  of 
tripotassium  phosphate,  and  suggests  as  a  cause  abnormal  exothermic 
heat  of  dissociation.  A  great  number  of  the  substances  studied  in  the 
present  investigation  showed  a  slight  decrease  in  dissociation  with  rise 
in  temperature.  Sodium  bromate,  sodium  thiosulphate,  trisodium 
phosphate,  and  lithium  chromate  showed  a  well-defined  increase. 
Several  compounds  have  almost  identical  dissociations  at  all  the  tem- 
peratures studied. 

TEMPERATURE  COEFFICIENTS. 

It  has  been  found  that  the  increase  in  conductivity  with  rise  in  tem- 
perature, is  due  primarily  to  the  velocities  with  which  the  ions  move. 
This  velocity  is  governed  by  the  viscosity  of  the  medium  and  the 
volume  and  mass  of  the  ion.  It  is  well  known  that  the  general  tendency 
of  rise  in  temperature  is  to  decrease  viscosity,  and  also  the  volume  and 
mass  of  the  ion,  if  the  ion  is  considered  not  as  a  charged  atom  or  group 
of  atoms,  but  as  a  charged  nucleus  plus  molecules  of  water,  which  must 
be  carried  along  in  all  migrations  through  the  remainder  of  the  solvent. 
Jones  has  given  a  number  of  proofs  for  the  validity  of  this  conception  of 
ions.  He  has  also  shown  that  these  complexes  break  down  at  higher 
temperatures.  With  these  facts  in  mind,  we  should  expect  a  greater 
increase  in  conductivity  with  rise  in  temperature  in  the  case  of  strongly 
hydrated  salts  than  in  the  case  of  weakly  hydrated  substances.  Taking 
the  amount  of  water  with  which  a  substance  crystallizes  as  indicative 
of  the  extent  to  which  it  is  hydrated,  it  is  found  that  all  of  the  com- 
pounds referred  to  in  the  above  tables  are  in  accord  with  this  concep- 
tion, except  potassium  ferricyanide  and  ammonium  chromate.  Ref- 
erence has  been  made  to  the  work  of  Jones  and  Getman  and  Jones  and 
Bassett,  which  throws  some  light  on  the  dissociation  of  the  complex 
ferricyanide,  and  also  shows  that  water  of  crystallization  is  not  always 
indicative  of  the  degree  of  hydration  to  which  a  compound  is  subject. 
While  they  proved  that  the  ferricyanide  is  not  hydrated,  it  is  rather 
probable  that  the  lithium  chromate  is,  since  lithium  salts,  as  a  class, 
have  a  much  greater  tendency  to  hydrate  than  the  salts  of  sodium  or 
potassium. 

The  temperature  coefficients,  expressed  in  percentage,  decrease  in  each 
case  with  rise  in  temperature.  They  increase  somewhat  on  dilution. 
This  is  especially  noticeable  with  hydrated  and  hydrolyzed  salts. 

'Amer.  Chem.  Journ.,  49,  249  (1913). 


160          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

Jones1  has  pointed  out  the  following  general  relations  deduced  from 
the  study  of  a  large  number  of  data  concerning  temperature  coefficients: 

1.  Those  ions  with  the  largest  hydrating  power  have  the  largest 
temperature  coefficients  of  conductivity. 

2.  Those  substances  having  equal  hydrating  power  have  approxi- 
mately the  same  temperature  coefficients  of  conductivity. 

3.  At  higher  dilutions,  the  temperature  coefficients  of  conductivity, 
for  any  given  substance,  are  greater  than  at  lower  dilutions. 

In  the  present  investigation  all  of  these  relations  have  been  found 
to  hold,  with  the  few  apparent  exceptions  noted  in  the  above  discussion. 

SUMMARY. 

1 .  A  brief  sketch  of  the  work  done  in  this  laboratory  on  the  conduc- 
tivity of  aqueous  solutions,  together  with  a  resume"  of  the  results 
obtained  by  Jones  and  his  co-workers,  has  been  given. 

2.  A  few  practical  suggestions  as  to  apparatus  and  methods  of  pro- 
cedure are  offered. 

3.  Eighteen  more  or  less  unusual  salts  were  studied  with  reference 
to  their  conductivity,  over  a  range  of  temperature  from  0°  to  35°. 
Their  temperature  coefficients  of  conductivity  are  expressed  in  two  sets 
of  units;  and  where  it  is  possible  their  dissociation  has  been  calculated. 

4.  The  results  of  this  investigation  are,  for  the  most  part,  in  accord 
with  the  findings  of  other  workers  in  this  field.     Three  exceptions  were 
found  to  the  rule  that  dissociation  decreases  with  rise  in  temperature. 
Two  apparent  exceptions  to  the  rule  that  large  temperature  coefficients 
of  conductivity  are  indicative  of  great  hydration  are  noted  and  possible 
explanations  are  offered. 

"Amer.  Chem.  Journ.,  34,  357  (1905). 


CHAPTER  IX. 
THE  DISSOCIATING  POWERS  OF  FREE  AND  OF  COMBINED  WATER. 


BY  G.  FRED.  OBDEMAN. 


INTRODUCTION. 

One  of  the  most  interesting  and  important  results  established  by 
Jones  and  Guy,1  in  their  work  with  the  radiomicrometer  on  the  absorp- 
tion spectra  of  solutions,  was  the  effect  of  the  dissolved  substance  on 
the  absorption  spectra  of  water.  They  noted  that  aqueous  solutions 
of  hydrated  salts  were  often  more  transparent  than  pure  water.  The 
absorption  of  aqueous  solutions  of  strongly  hydrated  salts  was  com- 
pared by  them  with  the  absorption  of  a  layer  of  water  equal  in  depth 
to  the  water  in  the  solution.  Similar  experiments  were  carried  out 
with  salts  which  were  but  slightly  hydrated.  The  concentrated  solu- 
tions of  strongly  hydrated  salts,  e.  g.,  5.3-normal  solution  of  calcium 
chloride,  were  found  to  be  more  transparent  than  a  comparable  quan- 
tity of  pure  water.  In  the  case  cited,  the  transparency  was  25  per  cent 
greater  from  1.05 /j,  to  1.2ju.  The  solutions  of  slightly  hydrated  salts, 
e.  g.,  potassium  chloride,  ammonium  chloride,  and  ammonium  nitrate, 
were  found  to  have,  in  general,  the  same  absorption  as  water  having 
the  same  depth  as  the  water  in  the  solution.  To  account  for  the  facts, 
they  concluded  that  the  combined  water  has  less  power  to  absorb  light 
than  free  or  uncombined  water.  This  is  regarded  by  them  as  striking 
evidence  that  some  of  the  water  in  the  presence  of  salts  which  by  other 
methods  are  shown  to  hydrate,  is  different  from  pure,  free,  uncombined 
water.  The  simplest  explanation  seems  to  be  that  this  is  combined 
water  or  water  of  hydration. 

The  work  of  Jones  and  Guy  has  been  repeated  and  extended  by 
Jones,  Shaeffer,  and  Paulus.2  Their  results  are  of  the  same  general 
character.  In  some  instances  they  found  the  aqueous  solutions  of 
hydrated  salts  to  be  40  per  .cent  more  transparent  than  a  comparable 
quantity  of  pure  water. 

Believing  a  determination  of  the  power  of  dissociation  of  combined 
water  might  aid  in  explaining  the  above  facts,  Dr.  Shaeffer  suggested 
the  use  of  isochloric  solutions  in  the  manner  we  have  adopted.  The 
object  of  this  work  has  been  to  ascertain  the  difference,  if  any,  between 
the  dissociating  power  of  combined  water  or  water  of  hydration,  and  the 
dissociating  power  of  uncombined  or  free  water. 

'Carnegie  Inst.  Wash.  Pub.  No.  190  (1913).    Phys.  Zeit.,  14, 278  (1913).        'Ibid.,  210  (1915). 

161 


162          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

EXPERIMENTAL  APPARATUS. 

The  Kohlrausch  method  of  measuring  conductivity  was  employed 
with  Wheatstone  bridge,  telephone  receiver,  and  induction  coil.  A  full 
description  of  the  apparatus  and  the  method  of  its  use  may  be  found  in 
the  earlier  publications  of  work  from  this  laboratory. 

Because  of  the  concentrated  solutions  employed,  a  different  type  of 
cell  was  necessary.  This  consisted  of  a  hard-glass,  U-shaped  tube, 
fitted  with  ground-glass  stoppers.  The  glass  tubes  carrying  the  elec- 
trodes were  sealed  by  means  of  sealing-wax  into  the  holes  bored  in  the 
centers  of  the  stoppers.  The  distance  between  the  electrodes  could 
thus  be  adjusted  if  occasion  demanded.  Numbers  were  etched  upon 
the  stoppers  and  the  corresponding  arms  of  the  U  -tubes,  so  that  the 
electrodes  could  always  be  placed  in  the  same  position. 

The  temperature  was  very  satisfactorily  maintained  by  means  of  a 
new  thermostat  described  in  detail  by  Jones,  Davis,  and  Putnam.1 

All  flasks  and  burettes  were  carefully  calibrated.  Solvent  and  solu- 
tions in  all  cases  were  brought  to  within  0.1°  of  the  necessary  tempera- 
ture before  measurements  were  made. 

SOLVENTS. 
WATER. 

The  water  was  purified  by  the  method  of  Jones  and  MacKay2 
as  modified  by  Schmidt,  and  had  a  mean  specific  conductivity  of 
1.5XlO-6at25°C. 

ISOCHLORIC  SOLUTIONS. 

Two  solutions  are  said  to  be  isochloric  if  they  contain  in  unit  volume 
the  same  number  of  chlorine  ions.  In  this  work  we  have  used  a  solu- 
tion of  potassium  chloride  normal  at  25°  ,  and  the  solution  of  calcium 
chloride  with  which  this  would  be  isochloric. 

Arrhenius3  shows,  for  the  condition  of  isohydric  solutions 

ma_n0 


Applying  this  to  the  solutions  under  discussion: 

a  =  percentage  dissociation  of  normal  solution  of  potassium  chloride; 

j8  =  that  of  the  corresponding  solution  of  calcium  chloride. 
Let  Vi  =  liters  of  solution  containing  a  gram-molecular  weight  of  potas- 
sium chloride  and  v2  =  corresponding  symbol  for  the  calcium  chloride. 
Let  m  =  number  of  chlorine  ions  per  molecule  of  potassium  chloride  and 
n  =  same  for  calcium  chloride. 

'Carnegie  Inst.  Wash.  Pub.  No.  210,  119  (1915).  3Zeit,  phys.  Chem.  2,  284  (1886). 

"Araer.  Chem.  Journ.,  17,  83  (1895). 


Dissociating  Powers  of  Free  and  of  Combined  Water.  163 

In  calculating  the  percentage  dissociations  by  the  conductivity 
method— 

*--*-  0--^  (2) 

Ml*,  MZoo 

Here  /ti  =  molecular  conductivity  of  the  solution  of  potassium  chloride, 
and  /z2  =  that  of  the  solution  of  calcium  chloride.  Further,  nix  =  con- 
ductivity at  infinite  dilution  for  the  potassium  chloride  and  p2oo  =  same 
for  calcium  chloride. 

From  the  method  of  Kohlrausch  for  calculating  conductivity, 


_  jf     "1  u\  _  jr     "2  ^2  /o\ 

1b1w1  2  6,  wt 

KI  and  K2  are  cell  constants.   01,  bi,  02,  and  &2  are  readings  on  the  bridge 
for  the  corresponding  resistances. 

Substituting  the  values  of  (3)  in  (2)  we  obtain 


I  t  t  2  2  8  ,  . 

a  =  ^—         -  and  ft  =  r  -  (4  ) 

biwlfilx  OtVtPi* 

Substituting  these  values  in  (1) 

mK1  Q!  vl         nKt  a2  vt 
*>i  wl  nlx  vl  ~  62  w2  KB  vt 

For  potassium  chloride  m  =  1  and  for  calcium  chloride  n  =  2.     If  the 
same  cell  be  used  then  Ki  =  K2.     And 


At  25°  MI  »  =  1371  and  ^2  x  =  246.5. 
Therefore, 


,     246.562 
Whence, 

0.8996  X         =  r 


Now,  by  measuring  a  and  6  of  the  solution  of  potassium  chloride  for 
the  resistance  wlt  the  left-hand  side  of  the  equation  becomes  a  constant. 
A  concentrated  solution  of  calcium  chloride  is  now  taken  in  different 

portions  and  diluted  in  different  amounts  until  ^-becomes  equal  to 

the  value  for  the  left-hand  side. 

The  conductivity  of  the  potassium  chloride  solution  was  found  to  be 
103.6.  The  specific  conductivity  of  the  calcium  chloride  solution  was 
found  to  be  93.  12.  The  calcium  chloride  solution  upon  analysis  proved 
to  be  0.6951  molar. 

'Carnegie  Inst.  Wash.  Pub.  No.  170,  p.  20  (1912). 


164          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

SALTS. 

The  sodium  chloride,  potassium  chloride,  and  ammonium  chloride 
were  the  best  grade  of  Baker's  analyzed  chemicals.  These  salts  were 
carefully  recrystallized  from  conductivity  water  and  thoroughly  dried. 
The  calcium  chloride,  magnesium  chloride,  and  strontium  chloride  were 
from  Eimer  and  Amend.  These  were  dissolved  in  conductivity  water 
and  filtered.  They  were  not  recrystallized,  being  so  soluble,  but,  as 
later  described,  were  used  as  concentrated  solutions. 

SOLUTIONS. 

Solutions  of  the  non-hydrated  salts  were  made  by  dissolving  the 
requisite  weights  of  salt  in  the  two  isochloric  solutions  and  diluting 
to  the  mark  with  these  two  solutions.  It  was  found  that  the  volume 
change  caused  by  the  added  salts  was  very  considerable  in  the  more 
concentrated  solutions.  This  means  that  the  solutions,  when  made, 
would  be  of  the  proper  strength  for  the  added  salts,  but  weaker  for  the 
isochloric  solutions.  Part  of  the  space  occupied  by  the  isochloric 
solutions  is  now  taken  by  the  added  salt.  However,  the  change  in 
volume  in  the  two  isochloric  solutions  was  the  same  for  the  same  added 
salt.  This  would  still  allow  the  results  to  be  comparable,  even  though 
no  correction  had  been  made  for  the  volume  change. 

Solutions  of  the  hydrated  salts  were  made  by  using  the  quantities 
of  a  concentrated  solution  of  known  strength  necessary  for  the  nor- 
mality desired.  Instead  of  using  isochloric  solutions,  the  necessary 
amounts  of  potassium  chloride  and  of  a  concentrated  solution  of  calcium 
chloride  were  employed.  The  solutions  were  now  brought  to  the  calibra- 
tion mark  by  conductivity  water.  In  this  way  a  solution  was  obtained 
accurate  with  respect  to  the  isochloric  solution  as  well  as  to  the  added 
salt.  The  strengths  of  the  various  concentrated  solutions  were  deter- 
mined by  an  estimation  of  the  chlorine  as  silver  chloride.  The  analyses 
were,  of  course,  made  in  duplicate.  All  solutions  were  made  up  at  25°  C. 
The  same  flask  was  used  for  normal  solutions  throughout.  This  was 
likewise  done  for  half-normal  and  eighth-normal  solutions. 

PROCEDURE. 

The  specific  conductivities  of  the  two  solutions  which  were  isochloric 
were  first  measured  and  computed  from  the  formula  s  =  K  —r*  The 

same  cell  was  employed  for  both  solutions,  so  that  any  change  in  the  cell 
constant  or  any  error  in  its  determination  would  be  eliminated  for 
comparison. 

Solutions  were  made  in  the  manner  described,  which  were  isochloric 
with  regard  to  the  potassium  chloride  and  calcium  chloride,  but  which 
were  also  of  a  known  normality  for  an  added  salt.  The  specific  con- 


Dissociating  Powers  of  Free  and  of  Combined  Water. 


165 


ductivities  of  these  solutions  were  now  determined.  These  were  not 
the  sum  of  the  specific  conductivities  of  the  two  salts  present  in  each 
case,  but  were  less,  because  of  the  common  ion  effect.  But  since  the 
two  solutions  contain  the  same  number  of  ions  (the  added  salt  not  being 
considered),  the  driving  back  of  the  dissociation  of  the  added  salt  by 
the  anions  would  be  the  same  in  each  case.  When  the  conductivity  of 
a  solution,  say,  normal  with  respect  to  potassium  chloride  and  half 
normal  with  respect  to  magnesium  chloride,  had  been  measured,  the 
cell  was  thoroughly  cleaned  and  dried.  This  same  cell  was  now  used 
for  the  determination  of  the  conductivity  of  a  solution  0.6951  normal 
with  respect  to  calcium  chloride,  that  is,  isochloric,  and  half-normal  as 
regards  magnesium  chloride.  Thus  the  possibility  of  error  due  to  any 
change  in  the  cell  was  eliminated. 

Three  concentrations  of  the  added  salt  were  used,  first  in  solutions 
normal  for  potassium  chloride  and  then  in  solutions  0.6957  normal  for 
calcium  chloride.  The  increase  in  conductivity  due  to  the  added  salt 
was  calculated  in  each  case.  The  difference  of  the  increases  for  com- 
parable solutions  was  then  computed. 

The  numbers  given  here  for  conductivities  represent  the  mean  of 
three  readings  of  the  bridge  for  different  resistances. 


TABLE  95. 


Normal  for  KC1. 


0.6951  normal  for  CaCl2. 


Added  salt. 

Specific 
conduc- 
tivity. 

Increase 
over  S  of 
mKCl. 

Specific 
conduc- 
tivity. 

Increase  over 
S  of  0.6951m 
CaCl,. 

Difference 
in 
increases. 

NaCl     8 

111.78 

8.18 

100.57 

7.45 

0.73 

NaCl     2 

135.78 

32.18 

120.33 

27.21 

4.97 

NaCl      1 

162.09 

58.49 

142.67 

49.55 

8.94 

KC1       8 

115.38 

11.78 

103.35 

10.23 

1.55 

KC1       2 

148.75 

45.15 

132.56 

39.44 

5.71 

KC1        1 

191.68 

88.08 

169.98 

76.86 

11.22 

NH4C1  8 

115.30 

11.70 

103.25 

10.13 

1.57 

NH<C1  2 

148.25 

44.65 

132.06 

38.84 

5.81 

NH4C1  1 

190.14 

86.54 

169.98 

76.86 

9.68 

MgCU   8 

117.02 

13.42 

104.03 

10.91 

2.51 

MgClj   2 

148.11 

44.51 

131.38 

38.26 

6.25 

MgCl2    1 

174.54 

70.94 

148.83 

55.71 

15.23 

CaCl2    8 

118.97 

15.37 

105.73 

12.61 

2.76 

Cads    2 

156.99 

53.39 

138.54 

45.42 

7.97 

CaCl2     1 

192.64 

89.04 

167.16 

74.04 

15.00 

SrCl2      8 

118.36 

14.70 

105.77 

12.65 

2.11 

SrClj      2 

156.28 

52.68 

137.22 

44.10 

8.58 

SrCl*      1 

192.35 

88.75 

167.29 

74.17 

14.58 

166          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

DISCUSSION  OF  RESULTS. 

The  figures  in  the  column  marked  "Increase  over  S  of  KC1"  mean 
the  increase  in  specific  conductivity  caused  by  the  amount  of  the  added 
salt  in  question,  over  the  specific  conductivity  due  entirely  to  the  potas- 
sium salt.  The  figures  in  column  marked  "Increase  over  S  of  0.6951m 
CaCl2"  mean  the  same  with  reference  to  the  solutions  of  calcium 
chloride,  isochloric  with  the  potassium  chloride.  The  last  column 
contains  the  difference  between  the  values  in  the  two  columns  just 
mentioned.  It  is  to  be  noticed,  in  all  the  cases  studied,  that  the 
increase  in  conductivity  is  greater  in  the  case  of  the  potassium  chloride. 
The  increase  for  the  last  three  added  salts  is  about  the  same.  What 
does  this  mean? 

The  driving  back  of  the  dissociation,  due  to  the  anions  of  the  iso- 
chloric solutions,  is  the  same  in  both  cases;  a  necessary  consequence  of 
the  solutions  being  isochloric.  Other  things  being  equal,  it  would  be 
expected  that  the  increase  in  conductivity  would  be  the  same  in  the 
two  solutions;  since  the  conductivity  is  a  measure  of  the  dissociation. 
How  can  the  above  facts  be  explained? 

We  have  a  tentative  explanation  which  is  by  no  means  final.  When 
a  salt  is  added  to  water  or  to  the  solution  of  another  salt,  the  added 
salt  is  dissociated  by  the  water  present.  We  believe  that  the  combined 
water,  in  the  solution  of  hydrated  salts,  is  less  associated  than  the  free 
water,  in  which  case  the  added  salts  would  be  less  dissociated;  since 
the  dissociation  power  of  a  solvent  changes  with  its  own  association. 
And  further,  this  effect  would  be  greater  the  greater  the  concentration, 
since  more  combined  water  would  then  be  present  in  the  solution  of 
calcium  chloride.  This  seems  to  be  in  accord  with  the  facts  established. 

These  results  and  conclusions  are  to  be  regarded  as  preliminary. 
Viscosity  undoubtedly  plays  a  role,  especially  in  the  more  concen- 
trated solutions.  The  atomic  volume  and  the  velocities  of  the  ions 
also  must  be  taken  into  account. 

Taking  all  of  the  factors  into  consideration,  these  results  render 
highly  probable  the  conclusion  that  the  dissociating  power  of  combined 
water  is  less  than  that  of  uncombined  water. 

In  conclusion,  we  would  thank  Dr.  E.  J.  Shaeffer,  who,  as  the  result 
of  his  spectroscopic  investigations  in  this  laboratory,  suggested  the 
study  of  this  problem. 


CHAPTER  X. 

THE  ABSORPTION  BY  SOILS  OF  POTASSIUM   FROM  AQUEOUS  SOLU- 
TIONS OF  POTASSIUM  CHLORIDE. 


BY  A.  G.  McCALL,  WITH  THE  COOPERATION  OP  F.  M.  HlLDEBRANDT,  F.  S.  HOLMES, 

E.  S.  JOHNSTON,  AND  S.  F.  TRELEASE. 


As  early  as  1866  Frank1  studied  the  retention  of  potassium  chloride 
by  the  soil,  using  metal  cylinders  3  inches  in  diameter  and  varying  in 
length  from  3  to  6  feet.  His  solution  contained  1  gram  of  potassium 
chloride  per  liter  of  water.  He  found  that  the  first  foot  of  soil  retained 
91  per  cent  of  the  potassium  chloride,  while  the  first  18  inches  removed 
95.5  per  cent  of  the  salt.  The  solution  appearing  at  the  bottom  of  his 
6-foot  columns  had  lost  all  but  2  per  cent  of  their  original  salt  content. 
The  addition  of  sodium  chloride  to  the  solution  diminished  the  absorp- 
tion of  the  potassium  chloride.2 

As  the  result  of  his  experiments  with  potassium  salts  Treutler3  con- 
cluded that  the  deeper  penetration  of  the  potassium  into  the  soil  was 
to  be  secured  by  the  use  of  potassium  chloride  rather  than  by  the  appli- 
cation of  potassium  sulphate  as  a  fertilizer. 

Peat  and  preparations  of  the  humic  acids  were  found  by  Heiden,4  to 
have  the  power  of  removing  a  part  of  the  salt,  when  brought  into  con- 
tact with  solutions  of  potassium  chloride.  The  greater  part  of  the 
salt  was  readily  recovered  by  the  use  of  a  small  quantity  of  water. 
Liebermann5  reported  that  aqueous  solutions  of  potassium  chloride 
showed  no  change  as  to  acidity  or  alkalinity  after  passing  through 
animal  charcoal,  but  that  the  concentration  was  decreased. 

More  recent  work  on  absorption  has  not  only  confirmed  these  earlier 
observations,  but  has  brought  out  the  fact  that  finely  divided  sub- 
stances exercise  a  selective  action  with  respect  to  the  solutions  with 
which  they  are  brought  in  contact.  In  some  cases  the  effect  of  this 
selective  action  is  to  remove  one  ion  of  the  salt  more  rapidly  than  the 
other,  leaving  the  solution  acid  or  alkaline,  depending  upon  which  ion 
is  absorbed  to  the  greater  extent.  Cameron  and  Bell6  found  that 
absorbent  cotton  has  the  power  of  removing  the  potassium  ion  from  a 
solution  of  potassium  chloride  more  rapidly  than  the  chlorine,  leaving 
the  solution  decidedly  acid  to  ordinary  indicators.  Previous  to  this 

"Landw.  Vers.-Stat.,  8,  45  (1866). 

'In  this  paper  the  terms  "absorption"  and  "adsorption"  are  used,  the  former  to  denote  the 
removal  of  material  from  solution  regardless  of  whether  the  act  is  essentially  chemical  or  physical, 
and  the  latter  to  designate  the  same  process  when  the  act  appears  to  be  clearly  physical,  that  is, 
dependent  upon  the  extent  of  surface  alone. 

"Landw.  Vers.-  Stat.,  12,  184  (1869);  15,  371  (1872). 

'Hoffman's  Jahresb.,  1866,  p.  29. 

'Wien.  Akad.  Ber.,  74,  331  (1877). 

•U.  8.  Dept.  Agriculture,  Bureau  of  Soils,  Bulletin  30  (1905). 

167 


168          Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 

work,  van  Bemmelen1  had  shown  that  the  treatment  of  a  soil  with  a 
solution  of  potassium  chloride  resulted  in  an  almost  complete  exchange 
of  potassium  for  sodium,  calcium,  and  magnesium.  In  one  experiment 
he  determined  the  chlorine  and  found  that  the  amount  had  not  changed. 

Schreiner  and  Failyer2  percolated  a  solution  of  potassium  chloride 
through  a  short  column  of  clay  soil  at  the  rate  of  50  c.c.  in  24  hours. 
The  first  few  hundred  cubic  centimeters  of  the  200  parts  per  million 
solution,  in  passing  through  the  soil,  was  reduced  in  concentration  to 
approximately  60  parts  per  million  of  potassium.  In  the  succeeding 
fractional  percolates  the  concentration  gradually  increased  until  it 
reached  173  parts  per  million  when  1,100  c.c.  had  passed.  Up  to  this 
point  the  soil  had  retained  approximately  900  parts  per  million  of 
potassium.  The  absorption  obtained  with  clay  loam  was  less  than 
that  observed  in  the  clay,  while  loam  soil  gave  results  intermediate 
between  the  clay  and  the  clay  loam.  For  the  sandy  loam  used,  the 
absorption  was  much  less  marked  than  in  the  finer  textured  soils,  but 
was,  nevertheless,  quite  marked  in  the  first  fractional  filtrates.  At 
the  close  of  the  absorption  periods  the  clay  and  the  clay  loam  soils  were 
washed  with  distilled  water,  the  percolation  of  water  being  at  the 
same  rate  as  that  used  in  passing  the  potassium  solution.  After  about 
450  c.c.  of  water  had  passed  through  the  clay  soil  the  successive  frac- 
tional percolates  showed  a  practically  constant  composition  of  20  parts 
per  1,000,000  of  potassium.  The  washing  was  continued  until  over 
2  liters  of  water  had  passed,  at  which  time  the  quantity  of  absorbed 
potassium  in  the  soil  had  been  reduced  to  350  parts  per  1,000,000  from 
an  initial  concentration  of  900  parts.  With  the  clay  loam  the  removal 
of  the  absorbed  salt  was  more  rapid.  The  quantity  of  potassium  in 
this  soil  was  reduced  from  570  to  250  parts  per  1,000,000  with  the 
passing  of  approximately  800  c.c.  of  water. 

Patton  and  Waggaman3  have  recalculated  the  results  of  Peters's4 
work  on  the  absorption  of  potassium  chloride  from  aqueous  solutions 
by  the  soil.  They  bring  out  the  fact  that  the  absorptive  effect  is  nearly 
twice  as  great  with  dilute  solutions  as  with  the  strong  concentrations. 
With  his  most  dilute  solution  (1.011  grams  of  potassium  per  kilo  of  soil) 
94  per  cent  of  the  total  potassium  present  was  absorbed,  while  from  a 
solution  20  times  as  strong  only  about  55  per  cent  was  removed. 

In  the  same  publication  Patton  and  Waggaman  call  attention  to  the 
fact  that  the  maximum  absorptive  capacity  of  an  absorbent,  while  a 
perfectly  definite  quantity,  is  of  little  practical  interest  in  soil  studies, 
because  of  the  fact  that  maximum  absorption  can  take  place  only  in 
the  presence  of  a  solution  which  is  saturated  with  respect  to  the  solute 
and  at  the  same  time  is  in  equilibrium  with  the  absorbing  material. 

'Landw.  Vers.-Stat.,  Bd.  21,  pp.  135-191  (1877). 

8U.  S.  Dept.  Agriculture,  Bureau  of  Soils,  Bui.  32  (1906). 

'/bid.,  Bui.  52,  (1908). 

'Landw.,  Vers.-Stat.,  2,  129  (1860). 


Absorption  by  Soils  of  Potassium.  169 

From  a  solution  which  is  less  than  saturated,  an  absorbent  can  remove 
a  quantity  of  the  solute  which  is  less  than  its  maximum  absorbent 
capacity,  but  which  is,  nevertheless,  a  definite  amount  for  any  particular 
concentration.  They  use  this  quantity  as  a  measure  of  the  specific 
absorptive  capacity  of  the  medium  with  respect  to  that  particular 
solution.  A  solution  of  potassium  chloride  brought  into  contact  with 
the  soil  will  lose  some  of  its  potassium  at  a  rate  which  gradually  de- 
creases until  the  salt  reaches  an  equilibrium  between  the  soil  and  the 
solution.  The  weight  of  potassium  chloride  absorbed  by  1  gram  of 
soil  represents  the  specific  absorptive  capacity  of  the  soil  for  that 
concentration  of  the  salt  at  that  particular  temperature.  If  conditions 
in  the  soil  are  altered,  more  salt  may  be  absorbed,  or  a  part  of  the  salt 
already  absorbed  may  go  back  into  solution.  Thus  it  is  seen  that  the 
absorption  process  is  a  reversible  phenomenon. 

In  summarizing  his  recent  work  at  the  Bureau  of  Soils,  Parker1  states: 

"The  rate  of  adsorption  of  the  chlorine  ions  from  a  solution  by  soils  is  much 
less  than  of  potassium  ions.  The  selective  adsorption  of  potassium  from  a 
potassium  chloride  solution  by  a  soil  increases  in  amount  with  the  concentra- 
tion up  to  a  certain  point,  and  then  remains  practically  constant.  In  general, 
the  smaller  the  soil  particles  the  greater  the  selective  adsorption  of  the  potas- 
sium from  a  potassium  chloride  solution  by  the  soil." 

Williams2  has  recently  called  attention  to  some  special  cases  of  selec- 
tive adsorption  to  which  he  has  given  the  name  "negative  adsorption." 
He  cites  the  work  of  Gore3  as  the  first  case  of  negative  adsorption. 
He  also  calls  attention  to  the  work  of  Lagergren,4  who  observed  that, 
upon  shaking  solutions  of  electrolytes  with  charcoal  or  silica,  the  con- 
centration of  the  salt  solution  increased  instead  of  decreasing.  Using 
blood  charcoal,  Williams  found  negative  adsorption  with  potassium 
chloride  at  certain  concentrations.  Up  to  a  concentration  of  0.0563 
gram  of  salt  per  gram  of  solution,  the  absorption  was  positive,  becom- 
ing negative  with  further  increase  in  concentration.  Gore's  work, 
cited  above,  gives  instances  of  negative  adsorption  at  low  concentra- 
tions, becoming  positive  in  less  dilute  solutions  of  the  same  salt.  In  all 
of  the  work  cited  the  investigators  have  had  their  interest  centered  upon 
the  condition  of  the  solution  after  equilibrium  had  been  reached;  conse- 
quently, the  solutions  were  left  in  contact  with  the  absorbing  medium 
for  from  24  hours  to  several  days.  In  the  extensive  studies  made  in 
the  Bureau  of  Soils,8  United  States  Department  of  Agriculture,  the 
contact  time  was  usually  a  24-hour  period. 

Throughout  the  work  in  which  soils  have  been  the  absorbing  medium, 
there  exists  a  very  great  deal  of  uncertainty  concerning  the  fundamental 
character  of  the  phenomenon.  The  evidence  from  the  use  of  chemi- 

'Journal  of  Agricultural  Research,  U.  S.  Dept.  Agriculture,  vol.  I,  No.  3  (1913). 

Trans.  Faraday  Society,  vol.  x,  part  1,  Aug.  1914. 

'Chemical  News,  69,  pp.  23,  33,  and  44  (1894). 

•Bihang  till  K.  Svenska  Vet-Akad.  Handlinger,  24,  n,  4  (1898). 

*Bul.  32  (1906)  and  Bui.  52  (1908). 


170 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


cally  inert  substances,  such  as  charcoal  and  silica,  would  indicate  that 
the  process  is  a  physical  one,  the  magnitude  of  the  adsorption  depending 
upon  the  extent  of  the  surface  presented  by  the  adsorbing  medium. 
On  the  other  hand,  some  of  the  very  earliest  work  with  soils  gave  strong 
evidence  of  the  chemical  replacement  of  the  bases  of  the  soil  by  the  base 
contained  in  the  solution.  With  these  facts  in  mind  a  series  of  experi- 
ments were  planned  for  the  purpose  of  studying,  (1)  the  effect  of  a  short 
time  contact  between  the  soil  and  the  salt  solution ;  and  (2)  the  effect 
of  the  extent  of  surface  upon  the  amount  and  the  rate  of  absorption. 

In  the  following  experiments  a  sample  of  Durham  sandy  loam1  soil  from 
Cabarrus  County,  North  Carolina,  was  used  as  the  absorbing  medium 
and  potassium  chloride  as  the  material  absorbed.  For  the  first  series  of 
experiments  a  sample  of  the  soil  was  dried  and  passed  through  a  2  mm. 
sieve.  For  the  second  series,  a  sample  of  the  same  soil  type  was 
reduced  to  very  fine  condition  by  grinding  for  4  days  in  a  porcelain- 
lined  ball  mill.  The  mechanical  analyses  of  the  two  samples,  as  made 
by  the  Bureau  of  Soils,  United  States  Department  of  Agriculture,  are 
given  in  table  96.  It  will  be  seen  from  the  table  that  practically  all 
of  the  sand  has  been  reduced  by  the  grinding  to  the  silt  and  clay 
groups,  thereby  making  an  enormous  increase  in  the  amount  of  surface 
exposed  by  the  two  samples  of  what  was  otherwise  identical  material. 
The  solution  used  was  potassium  chloride  of  approximately  five- 
hundredth  normal  concentration. 

TABLE  96. — Mechanical  analyses  of  soil  samples. 


Soil. 

Fine  gravel,  coarse 
sand,  medium  sand, 
2.0  to  0.25  mm. 

Fine  sand,  very 
fine  sand,  0.25 
to  0.05  mm. 

Silt,  0.05  to 
0.005  mm. 

Clay,  0.005 
to  0.0  mm. 

Durham  sandy  loam: 
Natural  soil  

47.1 

31.4 

17.9 

3.8 

Pulverized  four  days 
in  ball  mill 

0  5 

9  7 

78.1 

12.0 

DESCRIPTION  OF  THE  APPARATUS. 

In  order  to  secure  a  short  time  contact  between  the  soil  and  the  solu- 
tion a  special  apparatus  was  used,  whereby  the  solution  could  be  per- 
colated through  the  soil  at  any  desired  rate.  The  apparatus  consists 
of  a  porcelain-lined  filter  chamber  A  (fig.  17),  into  which  is  fitted  a 
short  filter  tube  B,  made  by  cutting  off  the  upper  end  of  a  Pasteur- 
Chamber  land  filter  tube  and  forcing  it  down  over  the  projection  on  the 
rubber  gasket  C.  Surrounding  the  filter  tube  is  a  brass  jacket  D, 
which  serves  as  a  container  for  the  soil.  After  the  introduction  of  the 
soil,  the  solution  is  poured  into  the  outer  jacket,  the  air-tight  cap  E  is 


'U.  S.  Bureau  of  Soils  classification,  Bui.  95,  p.  32. 


Absorption  by  Soils  of  Potassium. 


171 


screwed  down  over  the  top  of  the  filter  chamber,  and  the  system  con- 
nected to  an  automobile-tire  pump  through  the  valve  F.  The  filter 
tube  serves  to  hold  back  the  fine  soil  particles  and  give  a  clear  filtrate 
for  analysis.  The  apparatus  is  charged  by  putting  20  grams  of  the  air- 
dry  soil  in  the  brass  jacket  surrounding  the  filter  tube,  and  pouring  the 
salt  solution  into  the  porcelain-lined  chamber  which  surrounds  the  system. 
SERIES  1. 

Preliminary  to  the  use  of  the  soils  for  the 
absorption  work,  samples  of  both  the  natural 
and  the  finely  pulverized  soil  were  tested  for 
soluble  potassium,  by  percolating  pure  distilled 
water  through  them  at  the  approximate  rate  of 
200  c.c.  per  hour.  Each  successive  100  c.c.  por- 
tion of  the  filtrate  was  saved  and  separate  deter- 
minations made  for  potassium  by  the  calorimetric 
method  used  by  the  Bureau  of  Soils.1 

Table  97  contains  the  results  of  this  preliminary 
work.  The  numbers  in  the  first  column  indicate 
the  successive  filtrates,  while  the  second  column 
gives  the  concentrations  of  the  successive  portions 
of  the  filtrate  in  parts  per  million  of  the  solution. 
The  figures  of  the  table  will  be  used  in  the  subse- 
quent work,  in  order  to  correct  the  results  for  the 
soluble  potassium  originally  contained  in  the  soil. 
Although  it  has  no  direct  bearing  upon  the  ques- 
tions under  consideration,  it  is  of  interest  to  note 
that  the  grinding  of  the  soil  in  the  ball  mill  has 
increased  its  solubility  in  distilled  water  approximately  ten-fold. 

TABLE  97. — Water-soluble  potassium  in  samples  of  Durham  sandy  loam  soil. 


Parts  per  million  of  potassium 

Fractional 

in  filtrates. 

filtrates. 

Natural  soil. 

Pulverized  4  days 
in  ball  mill. 

1 

3.6 

37.5 

2 

2.2 

22.0 

3 

1.8 

20.0 

4 

1.6 

20.0 

SERIES  2. 

In  this  series  a  20-gram  sample  of  the  same  natural  soil  as  was  used 
in  series  1  was  percolated  with  250  c.c.  of  a  0.005  solution  of  potassium 
chloride,  the  flow  being  maintained  approximately  at  the  rate  of  50  c.c. 
in  10  minutes.  The  percolate  was  collected  in  fractions  of  50  c.c.  each, 

'U.  S.  Dept.  Agriculture,  Bureau  of  Soils,  Bui.  31,  p.  31. 


172 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


and  the  potassium  was  determined  calorimetrically.  Immediately 
following  the  percolation  with  the  solution  of  potassium  chloride,  pure 
distilled  water  was  forced  through  the  sample  at  the  same  rate  in  order 
to  study  the  removal  of  the  absorbed  potassium.  As  a  check  upon  the 
calorimetric  determinations,  the  specific  conductivity  of  the  fractional 
filtrates  was  determined  by  measuring  the  conductivity  of  the  solutions 
in  a  standardized  Arrhenius  cell,  according  to  the  method  described  by 
Jones.1 

TABLE  98. — Absorption  of  potassium  by  a  sandy  loam  soil,  from  a  solution  of  potassium  chloride 
containing  62  parts  per  million  of  potassium. 


Fractional 
filtrates. 

Parts  per  million 
of  potassium  in 
the  fractions. 

Potassium  retained 
in  parts  per  million 
of  dry  soil. 

Specific  con- 
ductivity 
at  0°. 

1 

40 

58 

92 

2 

36 

124 

82 

3 

40 

181 

79 

4 

44 

226 

79 

5 

59 

233 

78 

TABLE  99. — The  removal  of  the  absorbed  potassium  by  leaching  with  distilled  water. 


Fractional 
filtrates. 

Parta  per  million 
of  potassium  in 
the  fractions. 

Potassium  retained 
in  parts  per  million 
of  dry  soil. 

1 

7 

214 

2 

11 

186 

3 

9 

165 

4 

9 

142 

5 

9 

118 

6 

8 

99 

It  will  be  seen  from  table  98  that  the  first  10-minute  contact  of  the 
solution  with  the  soil  reduced  its  concentration  from  62  parts  per  million 
to  40  parts  per  million.  At  the  end  of  the  second  10  minute  period 
the  strength  of  the  solution  is  further  reduced  to  36  parts  per  million, 
but  from  this  point  the  concentration  of  the  solution  rises  until  the 
fifth  and  last  fraction  is  reached,  when  the  concentration  is  within  3 
parts  per  1,000,000  of  the  concentration  of  the  original  solution.  The 
third  column  shows  that  the  amount  of  potassium  retained  by  the  soil 
rises  gradually  to  233  parts  per  million  of  the  dry  soil,  when  250  c.c. 
of  solution  has  passed  through. 

Figures  18  and  19  represent  graphically  the  results  given  in  tables 
98  and  99.  In  figure  18  are  shown  the  results  based  upon  the  concen- 
tration of  the  successive  fractional  filtrates;  the  large  curve  representing 
the  concentration  in  parts  per  million  of  potassium,  and  the  small  curve 
the  specific  conductivities.  The  curve  for  the  potassium  shows  that 
the  first  three  fractions  of  the  filtrate  are  reduced  approximately  to 
two-thirds  of  the  strength  of  the  original  solution.  With  the  increase 


'Elements  of  Physical  Chemistry,  4th  edition,  pp.  377  to  383. 


Absorption  by  Soils  of  Potassium. 


173 


in  the  volume  of  the  filtrate  the  curve  rises  and  almost  reaches  the 
upper  boundary  of  the  figure,  which  line  represents  the  original  strength 
of  the  solution.  At  this  point  the  removal  of  the  absorbed  potassium 
is  begun  by  leaching  with  pure  distilled  water.  It  should  be  noted 
that  at  this  point  the  curve  falls  rapidly  to  practically  a  straight  line. 

Figure  19  gives  the  results  expressed  in  terms  of  the  soil,  the  abscissae 
being  the  volume  of  the  solution  or  water  percolated  through  the  soil, 
and  theordinates  the  amount  of  potassium  absorbed,  expressed  in  parts 
per  million  of  the  air-dry  soil.  While  it  is  possible,  and  indeed  quite 
probable  that  complete  equilibrium  was  not  reached,  it  is  apparent 


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Fio.  18. 

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Solution 


Water 


Fio.  19. 


that  the  soil  is  rapidly  approaching  a  saturated  condition  after  having 
absorbed  230  parts  per  1,000,000  of  potassium  as  a  result  of  the  perco- 
lation of  250  c.c.  of  the  salt  solution.  The  part  of  the  curve  repre- 
senting the  removal  of  the  absorbed  salt  by  leaching  with  distilled 
water  is  a  straight  line,  showing  the  uniformity  with  which  the  absorbed 
material  is  removed.  This  removal  of  the  absorbed  potassium  was 
accomplished  by  filling  the  filter  chamber  with  distilled  water  and 
continuing  the  percolation  at  the  same  rapid  rate  as  was  used  with  the 
solution  of  potassium  salt. 


174 


Conductivities  and  Viscosities  in  Pure  and  in  Mixed  Solvents. 


SERIES  3. 

For  series  3,  Durham  sandy  loam,  pulverized  four  days  in  a  ball  mill, 
was  used.  The  manipulation  was  the  same  as  for  series  2,  except  that 
the  fractional  percolates  were  100  c.c.  instead  of  50  as  in  the  previous 

TABLE  100. — Absorption  of  potassium  solution  by  a  finely  pulverized  sandy  loam  from  a  solution 
of  potassium  chloride  containing  78  parts  per  million  of  potassium. 


Part3  per  million  of  potassium  in  the  fractions. 

Specific  con- 

Fractional 

ductivity  of 

filtrates. 

Original 

After  correction  for 

solutions 

determinations. 

water-soluble  potassium. 

1 

144 

107 

112 

2 

112 

91 

82 

3 

100 

82 

82 

4 

112 

92 

84 

TABLE  101. — Removal  of  ab- 
sorbed potassium  by  leaching 
with  distilled  water. 


series,  making  the  total  volume  of  the  solution  twice  as  great.  An 
inspection  of  table  100  brings  out  the  surprising  fact  that  the  amount 
of  potassium  in  the  solution  has  been  increased  instead  of  decreased  by 
its  contact  with  the  soil. 

A  part  of  this  increase  in  concentration  is  undoubtedly  due  to  the 
fact  that  the  pulverized  soil  has  given  up  some  of  its  potassium  to  the 

percolating  solution.  However,  this  is  not 
sufficient  to  account  for  all  of  the  increase  in 
potassium  content.  In  column  3  of  table  100 
is  given  the  potassium  in  the  fractional  per- 
colates, after  they  have  been  corrected  for  the 
water-soluble  potassium  of  the  soil.  This  cor- 
rection is  based  upon  the  assumption  that 
pure  water  in  contact  with  the  soil  will  dis- 
solve potassium  at  the  same  rate,  and  in  the 
same  amounts  as  the  very  dilute  solution  of 
potassium  chloride  used  in  these  experiments. 
Proceeding  upon  this  assumption,  pure  water 

was  percolated  through  the  sample  of  pulverized  soil,  and  the  amount 
removed  by  the  water  was  applied  as  a  correction  to  the  correspond- 
ing fractional  percolates  of  the  salt  solution. 

The  results  given  in  table  100  are  shown  graphically  by  the  curves  in 
figure  21.  The  dotted  line  forming  the  upper  arm  of  the  curve  repre- 
sents the  original  concentration  of  the  successive  fractions,  while  the 
solid  line  represents  the  strength  of  the  solutions  after  the  correction 
has  been  applied.  It  will  be  seen  that,  even  after  the  correction  has 
been  made,  the  solution  still  maintains  a  concentration  higher  than  the 
original  salt  solution.  This  appears  to  be  a  case  of  selective  absorption, 
in  which  the  solvent  (water)  is  absorbed  more  rapidly  than  the  dissolved 


Fractional 
filtrates. 

Parts  per  million 
of  potassium  in 
the  fractions. 

1 

39 

2 

24 

3 

21 

4 

18 

5 

15 

6 

18 

Absorption  by  Soils  of  Potassium. 


175 


potassium  salt,  with  the  result  that  the  percolate  is  more  concentrated 
than  the  original  solution.  It  should  be  borne  in  mind,  however,  that 
we  are  probably  not  dealing  with  equilibrium  conditions.  The  mech- 
anism by  which  this  negative  absorption  is  effected,  may  be  explained 
by  assuming  that  the  solvent  and  the  dissolved  substance  are  capable 
of  being  absorbed  more  or  less  independently  and  at  different  rates. 
The  rapid  advance  of  the  liquid  through  the  fine  pores  of  the  soil  results, 
for  a  time,  in  the  more  rapid  absorption  of  the  water  than  of  the  salt, 


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Solution 


Fio.  21. 


leaving  the  liquid  in  the  larger  non-capillary  spaces  more  concentrated. 
This  more  concentrated  solution  then  moves  through  the  large  spaces 
and  appears  as  the  percolate.  It  is  quite  probable  that  after  a  few 
hundred  cubic  centimeters  have  passed  through,  equilibrium  will  have 
been  established  and  the  negative  absorption  will  become  positive. 
This  so-called  negative  absorption,  therefore,  may  be  considered  as 
a  special  case  of  selective  absorption. 

From  the  work  of  Williams,  to  which  reference  has  been  made,  it 
appears  that  the  absorption  of  the  solute  and  the  solvent  is  dependent 
upon  the  relative  masses  present.  With  some  electrolytes  in  water  it 
has  been  found  experimentally  that  the  absorption  is  at  first  positive, 
increases  to  a  maximum,  decreases  through  zero,  and  finally  becomes 
negative,  but  there  seems  to  be  no  case  on  record  in  which  the  initial 
absorption  effect  is  negative. 


Binder 

lord  Bros., 


Inc 


Makers 

Stockton,  Cam. 
f  H.  »N.  21.  1908 


9131 


Engineering 
Library 

THE  UNIVERSITY  OF  CALIFORNIA  LIBRARY 


